rs; 


AN   ELEMENTARY 


STUDY  OF  CHEMISTRY 


BY 


WILLIAM   McPHERSON,  PH.D. 

PROFESSOR  OF  CHEMISTRY,  OHIO  STATE  UNIVERSITY 


WILLIAM   EDWARDS  HENDERSON,  PH.D. 

ASSOCIATE  PROFESSOR  OF  CHEMISTRY,  OHIO  STATE  UNIVERSITY 


REVISED  EDITION 


GI.NN  &  COMPANY 

BOSTON  •  NEW  YORK  •  CHICAGO  •  LONDON 


COPYRIGHT,  1905,  1906,  BY 
WILLIAM  MCPHERSON  AND  WILLIAM  E.  HENDERSON 


ALL    RIGHTS    RESERVED 

68.1 


tgfte  g  then  gum 

GINN  &  COMPANY  •  PRO- 
PRIETORS •  BOSTON  •  U.S.A. 


PREFACE 

In  offering  this  book  to  teachers  of  elementary  chemistry 
the  authors  lay  no  claim  to  any  great  originality.  It  has 
been  their  aim  to  prepare  a  text-book  constructed  along 
lines  which  have  become  recognized  as  best  suited  to  an 
elementary  treatment  of  the  subject.  At  the  same  time 
they  have  made  a  consistent  effort  to  make  the  text  clear 
in  outline,  simple  in  style  and  language,  conservatively 
modern  in  point  of  view,  and  thoroughly  teachable. 

The  question  as  to  what  shall  be  included  in  an  elemen- 
tary text  on  chemistry  is  perhaps  the  most  perplexing  one 
which  an  author  must  answer.  While  an  enthusiastic 
chemist  with  a  broad  understanding  of  the  science  is  very 
apt  to  go  beyond  the  capacity  of  the  elementary  student, 
the  authors  of  this  text,  after  an  experience  of  many  years, 
cannot  help  believing  that  the  tendency  has  been  rather  in 
the  other  direction.  In  many  texts  no  mention  at  all  is  made 
of  fundamental  laws  of  chemical  action  because  their  com- 
plete presentation  is  quite  beyond  the  comprehension  of  the 
student,  whereas  in  many  cases  it  is  possible  to  present  the 
essential  features  of  these  laws  in  a  way  that  will  be  of 
real  assistance  in  the  understanding  of  the  science.  For 
example,  it  is  a  difficult  matter  to  deduce  the  law  of  mass 
action  in  any  very  simple  way  ;  yet  the  elementary  student 
can  readily  comprehend  that  reactions  are  reversible,  and 
that  the  point  of  equilibrium  depends  upon  rather  simple 
conditions.  The  authors  believe  that  it  is  worth  while  to 

iii 

221748 


iv  PREFACE 

present  such  principles  in  even  an  elementary  and  partial 
manner  because  they,  are  of  great  assistance  to  the  general 
student,  and  because  they  make  a  foundation  upon  which 
the  student  who  continues  his  studies  to  more  advanced 
courses  can  securely  build. 

The  authors  have  no  apologies  to  make  for  the  extent  to 
which  they  have  made  use  of  the  theory  of  electrolytic  dis- 
sociation. It  is  inevitable  that  in  any  rapidly  developing 
science  there  will  be  differences  of  opinion  in  regard  to  the 
value  of  certain  theories.  There  can  be  no  question,  how- 
ever, that  the  outline  of  the  theory  of  dissociation  here 
presented  is  in  accord  with  the  views  of  the  very  great 
majority  of  the  chemists  of  the  present  time.  Moreover, 
its  introduction  to  the  extent  to  which  the  authors  have 
presented  it  simplifies  rather  than  increases  the  difficulties 
with  which  the  development  of  the  principles  of  the  science 
is  attended. 

The  oxygen  standard  for  atomic  weights  has  been  adopted 
throughout  the  text.  The  International  Committee,  to 
which  is  assigned  the  duty  of  yearly  reporting  a  revised 
list  of  the  atomic  weights  of  the  elements,  has  adopted  this 
standard  for  their  report,  and  there  is  no  longer  any  author- 
ity for  the  older  hydrogen  standard.  The  authors  do  not 
believe  that  the  adoption  of  the  oxygen  standard  introduces 
any  real  difficulties  in  making  perfectly  clear  the  methods 
by  which  atomic  weights  are  calculated. 

The  problems  appended  to  the  various  chapters  have 
been  chosen  with  a  view  not  only  of  fixing  the  principles 
developed  in  the  text  in  the  mind  of  the  student,  but  also 
of  enabling  him  to  answer  such  questions  as  arise  in  his 
laboratory  work.  They  are,  therefore,  more  or  less  practi- 
cal in  character.  It  is  not  necessary  that  all  of  them  should 


PREFACE  V 

be  solved,  though  with  few  exceptions  the  lists  are  not 
long.  The  answers  to  the  questions  are  not  directly  given 
in  the  text  as  a  rule,  but  can  be  inferred  from  the  state- 
ments made.  They  therefore  require  independent  thought 
on  the  part  of  the  student. 

With  very  few  exceptions  only  such  experiments  are 
included  in  the  text  as  cannot  be  easily  carried  out  by  the 
student.  It  is  expected  that  these  will  be  performed  by 
the  teacher  at  the  lecture  table.  Directions  for  laboratory 
work  by  the  student  are  published  in  a  separate  volume. 

While  the  authors  believe  that  the  most  important  func- 
tion of  the  elementary  text  is  to  develop  the  principles  of 
the  science,  they  recognize  the  importance  of  some  discus- 
sion of  the  practical  application  of  these  principles  to  our 
everyday  life.  Considerable  space  is  therefore  devoted  to 
this  phase  of  chemistry.  The  teacher  should  supplement 
this  discussion  whenever  possible  by  having  the  class  visit 
different  factories  where  chemical  processes  are  employed. 

Although  this  text  is  now  for  the  first  time  offered  to 
teachers  of  elementary  chemistry,  it  has  nevertheless  been 
used  by  a  number  of  teachers  during  the  past  three  years. 
The  present  edition  has  been  largely  rewritten  in  the  light 
of  the  criticisms  offered,  and  we  desire  to  express  our 
thanks  to  the  many  teachers  who  have  helped  us  in  this 
respect,  especially  to  Dr.  William  Lloyd*  Evans  of  this  lab- 
oratory, a  teacher  of  wide  experience,  for  his  continued 
interest  and  helpfulness.  We  also  very  cordially  solicit 
correspondence  with  teachers  who  may  find  difficulties  or 

inaccuracies  in  the  text. 

THE  AUTHORS 

OHIO  STATE  UNIVERSITY,  COLUMBUS,  OHIO 
June,  1906 


CONTENTS 

CHAPTER  PAGE 

I.    INTRODUCTION i 

"II.   OXYGEN 13 

-III.   HYDROGEN 28 

-  IV.   WATER  AND  HYDROGEN  DIOXIDE 40 

V.   THE  ATOMIC  THEORY    .     .     .    4 59 

VI.   CHEMICAL  EQUATIONS  AND  CALCULATIONS      .     .  68 
-  VII.   NITROGEN    AND   THE   RARE   ELEMENTS   IN    THE 

ATMOSPHERE 78 

*"  VIII.   THE  ATMOSPHERE 83 

IX.   SOLUTIONS     T 94 

X.    ACIDS,  BASES,  AND  SALTS;  NEUTRALIZATION      .  106 

XL   VALENCE 116 

"~  XII.   COMPOUNDS  OF  NITROGEN 122 

XIII.    T"    /ERSIBLE    REACTIONS   AND    CHEMICAL    EQUI- 
LIBRIUM        137 

~~XIV.   SULPHUR  AND  ITS  COMPOUNDS 143 

""•x^XV.    PERIODIC  LAW  . 165 

XVI.   THE  CHLORINE  FAMILY      ..........  174 

XVII.    CARBON  AND  SOME  OF  ITS  SIMPLER  COMPOUNDS  196 

XVIII.    FLAMES,  —  ILLUMINANTS 213 

XIX.   MOLECULAR  WEIGHTS,  ATOMIC  WEIGHTS,  FORMULAS  223 

XX.   THE  PHOSPHORUS  FAMILY 238 

XXI.   SILICON,  TITANIUM,  BORON 257 

XXII.    THE  METALS 267 

XXIII.  THE  ALKALI  METALS     .     .     . 274 

XXIV.  THE  ALKALINE-EARTH  FAMILY .  300 

vii 


viii  CONTENTS 

CHAPTER  PAGE 

XXV.   THE  MAGNESIUM  FAMILY 316 

XXVI.   THE  ALUMINIUM  FAMILY 327 

XXVII.    THE  IRON  FAMILY 338 

XXVIII.   COPPER,  MERCURY,  AND  SILVER 356 

XXIX.   TIN  AND  LEAD 370 

XXX.   MANGANESE  AND  CHROMIUM 379 

XXXI.   GOLD  AND  THE  PLATINUM  FAMILY 390 

XXXII.   SOME  SIMPLE  ORGANIC  COMPOUNDS 397 

INDEX 421 

APPENDIX  A Facing  back  cover 

APPENDIX  B  .     .     Inside  back  cover 


AN  ELEMENTARY   STUDY 
OF   CHEMISTRY 


CHAPTER  I 
INTRODUCTION 

The  natural  sciences.  Before  we  advance  very  far  in  the 
study  of  nature,  it  becomes  evident  that  the  one  large  study 
must  be  divided  into  a  number  of  more  limited  ones  for  the 
convenience  of  the  investigator  as  well  as  of  the  student. 
These  more  limited  studies  are  called  the  natural  sciences. 

Since  the  study  of  nature  is  divided  in  this  way  for  mere 
convenience,  and  not  because  there  is  any  division, in  nature 
itself,  it  often  happens  that  the  different  sciences  are  very 
intimately  related,  and  a  thorough  knowledge  of  any  one 
of  them  involves  a  considerable  acquaintance  with  several 
others.  Thus  the  botanist  must  know  something  about 
animals  as  well  as  about  plants ;  the  student  of  human 
physiology  must  know  something  about  'physics  as  well  as 
about  the  parts  of  the  body. 

Intimate  relation  of  chemistry  and  physics.  Physics 
and  chemistry  are  two  sciences  related  in  this  close  way, 
and  it  is  not  easy  to  make  a  precise  distinction  between 
them.  In  a  general  way  it  may  be  said  that  they  are  both 
concerned  with  inanimate  matter  rather  than  with  living, 
and  more  particularly  with  the  changes  which  such  matter 


2          AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

may  be  made  to  undergo.  These  changes  must  be  consid- 
ered more  closely  before  a  definition  of  the  two  sciences 
can  be  given. 

Physical  changes.  One  class  of  changes  is  not  accompa- 
nied by  an  alteration  in  the  composition  of  matter.  When 
a  lump  of  coal  is  broken  the  pieces  do  not  differ  from  the 
original  lump  save  in  size.  A  rod  of  iron  may  be  broken 
into  pieces ;  it  may  be  magnetized  ;  it  may  be  heated  until 
it  glows ;  it  may  be  melted.  In  none  of  these  changes  has 
the  composition  of  the  iron  been  affected.  The  pieces  of 
iron,  the  magnetized  iron,  the  glowing  iron,  the  melted  iron, 
are  just  as  truly  iron  as  was  the  original  rod.  Sugar  may 
be  dissolved  in  water,  but  neither  the  sugar  nor  the  water  is 
changed  in  composition.  The  resulting  liquid  has  the  sweet 
taste  of  sugar ;  moreover  the  water  may  be  evaporated  by 
heating  and  the  sugar  recovered  unchanged.  Such  changes 
are  called  physical  changes. 

DEFINITION  :  Physical  changes  are  those  which  do  not 
involve  a  change  in  the  composition  of  the  matter. 

Chemical  changes.  Matter  may  undergo  other  changes 
in  which  its  composition  is  altered.  When  a  lump  of  coal 
is  burned  ashes  and  invisible  gases  are  formed  which  are 
entirely  different  in  composition  and  properties  from  the 
original  coal.  A  rod  of  iron  when  exposed  to  moist  air  is 
gradually  changed  into  rust,  which  is  entirely  different  from 
the  original  iron.  When  sugar  is  heated  a  black  substance 
is  formed  which  is  neither  sweet  nor  soluble  in  water.  Such 
changes  are  evidently  quite  different  from  the  physical 
changes  just  described,  for  in  them  new  substances  are 
formed  in  place  of  the  ones  undergoing  change.  Changes 
of  this  kind  are  called  chemical  changes. 


INTRODUCTION  3 

DEFINITION  :  Chemical  changes  are  those  which  involve  a 
change  in  the  composition  of  the  matter. 

How  to  distinguish  between  physical  and  chemical  changes.  It  is 
not  always  easy  to  tell  to  which  class  a  given  change  belongs,  and 
many  cases  will  require  careful  thought  on  the  part  of  the  student. 
The  test  question  in  all  cases  is,  Has  the  composition  of  the  sub- 
stance been  changed  ?  Usually  this  can  be  answered  by  a  study  of 
the  properties  of  the  substance  before  and  after  the  change,  since  a 
change  in  composition  is  attended  by  a  change  in  properties.  In 
some  cases,  however,  only  a  trained  observer  can  decide  the  question. 

Changes  in  physical  state.  One  class  of  physical  changes 
should  be  noted  with  especial  care,  since  it  is  likely  to  prove 
misleading.  It  is  a  familiar  fact  chat  ice  is  changed  into 
water,  and  water  into  steam,  by  heating.  Here  we  have 
three  different  substances,  —  the  solid  ice,  the  liquid  water, 
and  the  gaseous  steam,  —  the  properties  of  which  differ 
widely.  The  chemist  can  readily  show,  however,  that  these 
three  bodies  have  exactly  the  same  composition,  being  com- 
posed of  the  same  substances  in  the  same  proportion.  Hence 
the  change  from  one  of  these  substances  into  another  is  a 
physical  change.  Many  other  substances  may,  under  suit- 
able conditions,  be  changed  from  solids  into  liquids,  or  from 
liquids  into  gases,  without  change  in  composition.  Thus 
butter  and  wax  will  melt  when  heated  ;  alcohol  and  gasoline 
will  evaporate  when  exposed  to  the  air.  "The  three  states  — 
solid,  liquid,  and  gas  —  are  called  the  three  physical  states 
of  matter. 

Physical  and  chemical  properties.  Many  properties  of  a 
substance  can  be  noted  without  causing  the  substance  to 
undergo  chemical  change,  and  are  therefore  called  its  phys- 
ical properties.  Among  these  are  its  physical  statcy  -eelor, 
odor,  taste,  size,  shape,  weight.  Other  properties  are  only 


4         AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

discovered  when  the  substance  undergoes  chemical  change. 
These  are  called  its  chemical  properties.  Thus  we  find  that 
coal  burns  in  air,  gunpowder  explodes  when  ignited,  milk 
sours  when  exposed  to  air. 

Definition  of  physics  and  chemistry.  It  is  now  possible  to 
make  a  general  distinction  between  physics  and  chemistry. 

DEFINITION  :  Physics  is  the  science  which  deals  with 
those  changes  in  matter  which  do  not  involve  a  change  in 
composition. 

DEFINITION  :  Chemistry  is  the  science  which  deals  with 
those  changes  in  matter  which  do  involve  a  change  in  com- 
position. 

Two  factors  in  all  changes.  In  all  the  changes  which 
matter  can  undergo,  whether  physical  or  chemical,  two 
factors  must  be  taken  into  account,  namely,  energy  and 
matter. 

Energy.  It  is  a  familiar  fact  that  certain  bodies  have 
the  power  to  do  work.  Thus  water  falling  from  a  height 
upon  a  water  wheel  turns  the  wheel  and  in  this  way  does 
the  work  of  the  mills.  Magnetized  iron  attracts  iron  to 
itself  and  the  motion  of  the  iron  as  it  moves  towards  the 
magnet  can  be  made  to  do  work.  When  coal  is  burned  it 
causes  the  engine  to  move  and  transports  the  loaded  cars 
from  place  to  place.  When  a  body  has  this  power  to  do  work 
it  is  said  to  possess  energy. 

Law  of  conservation  of  energy.  Careful  experiments 
have  shown  that  when  one  body  parts  with  its  energy  the 
energy  is  not  destroyed  but  is  transferred  to  another  body 
or  system  of  bodies.  Just  as  energy  cannot  be  destroyed, 
neither  can  it  be  created.  If  one  body  gains  a  certain  amount 
of  energy,  some  other  body  has  lost  an  equivalent  amount 


INTRODUCTION  5 

These  facts  are  summed  up  in  the  law  of  conservation  of 
energy  which  may  be  stated  thus:  While  energy  can  be 
changed  from  one  form  into  another y  it  cannot  be  created  or 
destroyed. 

Transformations  of  energy.  Although  energy  can  neither 
be  created  nor  destroyed,  it  is  evident  that  it  may  assume 
many  different  forms.  Thus  the  falling  water  may  turn  the 
electric  generator  and  produce  a  current  of  electricity.  The 
energy  lost  by  the  falling  water  is  thus  transformed  into 
the  energy  of  the  electric  current.  This  in  turn  may  be 
changed  into  the  energy  of  motion,  as  when  the  current  is 
used  for  propelling  the  cars,  or  into  the  energy  of  heat 
and  light,  as  when  it  is  used  for  heating  and  lighting  the 
cars.  Again,  the  energy  of  coal  may  be  converted  into 
energy  of  heat  and  subsequently  of  motion,  as  when  it  is 
used  as  a  fuel  in  steam  engines. 

Since  the  energy  possessed  by  coal  only  becomes  avail- 
able when  the  coal  is  made  to  undergo  a  chemical  change, 
it  is  sometimes  called  chemical  energy.  It  is  this  form  of 
energy  in  which  we  are  especially  interested  in  the  study 
of  chemistry. 

Matter.  Matter  may  be  defined  as  that  which  occu- 
pies space  and  possesses  weight.  Like  energy,  matter  may 
be  changed  oftentimes  from  one  form  into  another ;  and 
since  in  these  transformations  all  the  other  physical  proper- 
ties of  a  substance  save  weight  are  likely  to  change,  the 
inquiry  arises,  Does  the  weight  also  change  ?  Much  careful 
experimenting  has  shown  that  it  does  not.  The  weight  of 
the  products  formed  in  any  change  in  matter  always  equals 
the  weight  of  the  substances  undergoing  change. 

Law  of  conservation  of  matter.  The  important  truth  just 
stated  is  frequently  referred  to  as  the  law  of  conservation 


AN   ELEMENTARY  STUDY  OF  CHEMISTRY 


of  matter,  and  this  law  may  be  briefly  stated  thus  :  Matter 
can  neither  be  created  nor  destroyed,  though  it  can  be  changed 
from  one  form  into  another. 

Classification  of  matter.  At  first  sight  there  appears  to 
be  no  limit  to  the  varieties  of  matter  of  which  the  world  is 
made.  For  convenience  in  study  we  may  classify  all  these 
varieties  under  three  heads,  namely,  mechanical  mixtures, 
chemical  compounds,  and  elements. 

Mechanical  mixtures.  If  equal  bulks  of  common  salt  and 
iron  filings  are  thoroughly  mixed  together,  a  product  is 
obtained  which,  judging  by  its  appearance,  is  a 
new  substance.  If  it  is  examined  more  closely, 
however,  it  will  be  seen  to  be  merely  a  mixture 
of  the  salt  and  iron,  each  of 
which  substances  retains  its  own 
peculiar  properties.  The  mixture 
tastes  just  like  salt ;  the  iron 
particles  can  be  seen  and  their 
gritty  character  detected.  A 
magnet  rubbed  in  the  mixture 
draws  out  the  iron  just  as  if  the 
salt  were  not  there.  On  the  other 
hand,  the  salt  can  be  separated 
from  the  iron  quite  easily.  Thus,  if  several  grams  of  the 
mixture  are  placed  in  a  test  tube,  and  the  tube  half  filled 
with  water  and  thoroughly  shaken,  the  salt  dissolves  in  the 
water.  The  iron  particles  can  then  be  filtered  from  the 
liquid  by  pouring  the  entire  mixture  upon  a  piece  of  filter 
paper  folded  so  as  to  fit  into  the  interior  of  a  funnel  (Fig.  i). 
The  paper  retains  the  solid  but  allows  the  clear  liquid, 
known  as  the  filtrate,  to  drain  through.  The  iron  particles 
left  upon  the  filter  paper  will  be  found  to  be  identical  with 


FIG.  i 


c= 


INTRODUCTION  7 

the  original  iron.  The  salt  can  be  recovered  from  the  ni- 
trate by  evaporation  of  the  water.  To  accomplish  this  the 
nitrate  is  poured  into  a  small  evaporating  dish  and  gently 
heated  (Fig.  2)  until  the  water  has  disappeared,  or  evapo- 
rated. The  solid  left  in  the  dish  is  identical  in  every  way 
with  the  original  salt.  Both  the  iron  and  the  salt  have  thus 
been  recovered  in  their  original  condition.  It  is  evident  that 
no  new  substance  has  been  formed  by  rubbing  the 
salt  and  iron  together.  The  product  is  called  a 
mechanical  mixture.  Such  mixtures  are  very  com- 
mon in  nature,  almost  all  minerals,  sands,  and  soils 
being  examples  of  this  class  of  substances.  It  is  at 
once  apparent  that  there  is  no  law  regulating  the 
composition  of  a  mechanical  mixture, 
and  no  two  mixtures  are  likely  to  have 
exactly  the  same  composition.  The 
ingredients  of  a  mechanical  mixture 
can  usually  be  separated  by  mechani- 
-  cal  means,  such  as  sifting,  sorting, 
magnetic  attraction,  or  by  dissolving 
one  constituent  and  leaving  the  other  unchanged. 

DEFINITION  :  A  mechanical  mixture  is  one  in  which  the 
constituents  retain  their  original  properties,  no  chemical  action 
having  taken  place  when  they  were  brought  together. 

Chemical  compounds.  If  iron  filings  and  powdered  sul- 
phur are  thoroughly  ground  together  in  a  mortar,  a  yellow- 
ish-green substance  results.  It  might  easily  be  taken  to  be 
a  new  body ;  but  as  in  the  case  of  the  iron  and  salt,  the  in- 
gredients can  readily  be  separated.  A  magnet  draws  out  the 
iron.  Water  does  not  dissolve  the  sulphur,  but  other  liquids 
do,  as,  for  example,  the  liquid  called  carbon  disulphide. 


8 


AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


When  the  mixture  is  treated  with  carbon  disulphide  the  iron 
is  left  unchanged,  and  the  sulphur  can  be  obtained  again, 
after  filtering  off  the  iron,  by  evaporating  the  liquid.  The 
substance  is,  therefore,  a  mechanical  mixture. 

If  now  a  new  portion  of  the  mixture  is  placed  in  a  dry 
test  tube  and  carefully  heated  in  the  flame  of  a  Bunsen 
burner,  as  shown  in  Fig.  3,  a  striking  change  takes  place. 
The  mixture  begins  to  glow  at  some 
point,  the  glow  rapidly  extending 
throughout  the  whole  mass.  If  the 
test  tube  is  now  broken  and  the  prod- 
uct examined,  it  will  be  found  to  be 
a  hard,  black,  brittle  substance,  in 
no  way  recalling  the  iron  or  the 
sulphur.  The  magnet  no  longer  at- 
tracts sit ;  carbon  disulphide  will  not 
dissolve  sulphur  from  it.  It  is  a  new 
substance  with  new  properties,  resulting  from  the  chemical 
union  of  iron  and  sulphur,  and  is  called  iron  sulphide.  Such 
substances  are  called  chemical  compounds,  and  differ  from 
mechanical  mixtures  in  that  the  substances  producing  them 
lose  their  own  characteristic  properties.  We  shall  see  later 
that  the  two  also  differ  in  that  the  composition  of  a  chemical 
compound  never  varies. 

DEFINITION  :  A  chemical  compound  is  a  substance  the  con- 
stituents of  which  have  lost  their  own  charade*  istic properties, 
and  which  cannot  be  separated  save  by  a  chemical  change. 

Elements.  It  has  been  seen  that  iron  sulphide  is  com- 
posed of  two  entirely  different  substances,  — iron  and  sulphur. 
The  question  arises,  Do  these  substances  in  turn  contain 
other  substances,  that  is,  are  they  also  chemical  compounds  ? 


FIG.  3 


INTRODUCTION  9 

Chemists  have  tried  in  a  great  many  ways  to  decompose 
them,  but  all  their  efforts  have  failed.  Substances  which 
have  resisted  all  efforts  to  decompose  them  into  other  sub- 
stances are  called  elements.  It  is  not  always  easy  to  prove 
that  a  given  substance  is  really  an  element.  Some  way  as 
yet  untried  may  be  successful  in  decomposing  it  into  other 
simpler  forms  of  matter,  and  the  supposed  element  will 
then  prove  to  be  a  compound.  Water,  lime,  and  many 
other  familiar  compounds  were  at  one  time  thought  to 
be  elements. 

DEFINITION  :  An  element  is  a  substance  which  cannot  be 
separated  into  simpler  substances  by  any  known  means. 

Kinds  of  matter.  While  matter  has  been  grouped  in  three  classes 
for  the  purpose  of  study,  it  will  be  apparent  that  there  are  really 
but  two  distinct  kinds  of  matter,  namely,  compounds  and  elements. 
A  mechanical  mixture  is  not  a  third  distinct  kind  of  matter,  but  is 
made  up  of  varying  quantities  of  either  compounds  or  elements  or 
both. 

Alchemy.  In  olden  times  it  was  thought  that  some  way  could  be 
found  to  change  one  element  into  another,  and  a  great  many  efforts  were 
made  to  accomplish  this  transformation.  Most  of  these  efforts  were 
directed  toward  changing  the  commoner  metals  into  gold,  and  many 
fanciful  ways  for  doing  this  were  described.  The  chemists  of  that 
time  were  called  alchemists,  and  the  art  which  they  practiced  was 
called  alchemy.  The  alchemists  gradually  became  convinced  that  the 
only  way  common  metals  could  be  changed  into  gold  was  by  the  won- 
derful power  of  a  magic  substance  which  they  called  the  philosopher's 
stone,  which  would  accomplish  this  transformation  by  its  mere  touch 
and  would  in  addition  give  perpetual  youth  to  its  fortunate  possessor. 
No  one  has  ever  found  such  a  stone,  and  no  one  has  succeeded  in 
changing  one  metal  into  another. 

e  number  of  substances  now  con- 
out  eighty  in  all. 
make  any 


10       AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

large  fraction  of  the  materials  in  the  earth's  crust.  Clarke 
gives  the  following  estimate  of  the  composition  of  the 
earth's  crust : 

Oxygen 47-Q%      Calcium 3.5% 

Silicon 27-9         Magnesium 2.5 

Aluminium     .....       8.1         Sodium 2.7 

Iron 4.7         Potassium 2.4 

Other  elements  ....     1.2% 

A  complete  list  of  the  elements  is  given  in  the  Appendix. 
In  this  list  the  more  common  of  the  elements  are  marked 
with  an  asterisk.  It  is  not  necessary  to  study  more  than 
a  third  of  the  total  number  of  elements  to  gain  a  very  good 
knowledge  of  chemistry. 

Physical  state  of  the  elements.  About  ten  of  the  elements 
are  gases  at  ordinary  temperatures.  Two  —  mercury  and 
bromine  —  are  liquids.  The  others  are  all  solids,  though 
their  melting  points  vary  through  wide  limits,  from  caesium 
which  melts  at  26°  to  elements  which  do  not  melt  save  in 
the  intense  heat  of  the  electric  furnace. 

Occurrence  of  the  elements.  Comparatively  few  of  the 
elements  occur  as  uncombined  substances  in  nature,  most 
of  them  being  found  in  the  form  of  chemical  compounds. 
When  an  element  does  occur  by  itself,  as  is  the  case  with 
gold,  we  say  that  it  occurs  in  the  free  state  or  native ;  when 
it  is  combined  with  other  substances  in  the  form  of  com- 
pounds, we  say  that  it  occurs  in  the  combined  state,  or  in 
combination.  In  the  latter  case  there  is  usually  little  about 
the  compound  to  suggest  that  the  element  is  present  in  it ; 
for  we  have  seen  that  elements  lose  their  own  peculiar  proper- 
ties when  they  enter  into  combination  with  other  elements. 
It  would  never  be  suspected,  for  example,  that  the  reddish, 
earthy-looking  iron  ore  contains  iron. 


INTRODUCTION  1 1 

Names  of  elements.  The  names  given  to  the  elements 
have  been  selected  in  a  great  many  different  ways.  ( i )  Some 
names  are  very  old  and  their  original  meaning  is  obscure. 
Such  names  are  iron,  gold,  and  copper.  (2)  Many  names  in- 
dicate some  striking  physical  property  of  the  element.  The 
name  bromine,  for  example,  is  derived  from  a  Greek  word 
meaning  a  stench,  referring  to  the  extremely  unpleasant 
odor  of  the  substance.  The  name  iodine  comes  from  a  word 
meaning  violet,  alluding  to  the  beautiful  color  of  iodine  vapor. 
(3)  Some  names  indicate  prominent  chemical  properties  of 
the  elements.  Thus,  nitrogen  means  the  producer  of  niter, 
nitrogen  being  a  constituent  of  niter  or  saltpeter.  Hydro- 
gen means  water .  former,  signifying  its  presence  in  water. 
Argon  means  lazy  or  inert,  the  element  being  so  named 
because  of  its  inactivity.  (4)  Other  elements  are  named 
from  countries^  or  localities,  as  germanium  and  scandium. 

Symbols.  In  indicating  the  elements  found  in  compounds 
it  is  inconvenient  to  use  such  long  names,  and  hence  chem- 
ists have  adop$e$  a  system  of  abbreviations.  These  abbre- 
viations area  town  as  symbols,  each  element  having  a 
distinctive  symBol.  (i)  Sometimes  the  initial  letter  of  the 
name  will  ^uj|9  to  indicate  the  element.  Thus  I  stands 
for  iodine,  C  ferr  carbon.  (2)  Usually  it  is  necessary  to  add 
some  other  characteristic  letter  to  the  symbol,  since  several 
names  may  be'giii  with  the  same  letter.  Thus  C  stands  for 
carbon,  Cl  for  chlorine,  Cd  for  cadmium,  Ce  for  cerium,  Cb 
for  coluir»biiir~  (3)  Sometimes  the  symbol  is  an  abbrevia- 
tion of  th<  Latin  name.  In  this  way  Fe  (ferrum)  indi- 
cates :  Cuprum),  copper,  Au  (aurum),  gold.  The 
symbols  included  in  the  list  of  elements  given  in  the 
Appendi  y  will  become  familiar  through  constant 
use. 


12       AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Chemical  affinity  the  cause  of  chemical  combination.   The 

agency  which  causes  substances  to  combine  and  which  holds 
them  together  when  combined  is  called  chemical  affinity. 
The  experiments  described  in  this  chapter,  however,  show 
that  heat  is  often  necessary  to  bring  about  chemical  action. 
The  distinction  between  the  cause  producing  chemical  action 
and  the  circumstances  favoring  it  must  be  clearly  made. 
Chemical  affinity  is  always  the  cause  of  chemical  union. 
Many  agencies  may  make  it  possible  for  chemical  affinity 
to  act  by  overcoming  circumstances  which  stand  in  its  way. 
Among  these  agencies  are  heat,  light,  and  electricity.  As 
a  rule,  solution  also  promotes  action  between  two  substances. 
Sometimes  these  agencies  may  overcome  chemical  .attrac- 
tion and  so  occasion  the  decomposition  of  a  compound. 

EXERCISES 

1.  To  what  class  of  changes  do  the  following  belong?    (a)  The 
melting  of  ice  ;   (b}  the  souring  of  milk  ;   (c)  the  burning  of  a  candle  ; 
(W)  the  explosion  of  gunpowder ;  (e}  the  corrosion  of  metals.    What 
test  question  must  be  applied  in  each  of  the  above  cases  ? 

2.  Give  two  additional  examples  (a)  of  chemical  changes ;  (£)  of 
physical  changes. 

3.  Is  a  chemical  change  always  accompanied  bj^piysical  change  ? 
Is  a  physical  change  always  accompanied  by  a  chemical  change  ? 

4.  Give  two  or  more  characteristics  of  a  chemical  change. 

5.  (a)  When  a  given  weight  of  water  freezes,  does  it  absorb  or 
evolve  heat  ?  (£)  When  the  resulting  ice  melts,  is  the  total  heat  change 
the  same  or  different  from  that  of  freezing  ? 

6.  Give  three  examples  of  each  of  the  following :  (#)  mechanical 
mixtures;  (£)  chemical  compounds;'  (<;)  elements. 

7.  Give  the  derivation  of  the  names  of  the  following  elements: 
thorium,  gallium,  selenium,  uranium.      (Consult  dictionary.) 

8.  Give  examples  of  chemical  changes  which  are  produced  through 
the  agency  of  heat ;  of  light ;  of  electricity. 


CHAPTER    II 

^ 

OXYGEN 

History.  The  discovery  of  oxygen  is  generally  attributed 
to  the  English  chemist  Priestley,  who  in  1 774  obtained  the 
element  by  heating  a  compound  of  mercury  and  oxygen, 
known  as  red  oxide  of  mercury.  It  is  probable,  however, 
that  the  Swedish  chemist  Scheele  had  previously  obtained  it, 
although  an  account  of  his  experiments  was  not  published 
until  1777-  The  name  oxygen  signifies  acid  former.  It 
was  given  to  the  element  by  the  French  chemist  Lavoisier, 
since  he  believed  that  all  acids  owe  their  characteristic 
properties  to  the  presence  of  oxygen.  This  view  we  now 
know  to  be  incorrect. 

Occurrence.    Oxygen  is  by  far  the  most  abundant  of  all 
the  elements.    It  occurs  both  in  the  free  and  in  the  com- 
bined state.    In  the  free  state  it  occurs  in  the  air,  100  vol-     j 
umes  of  dry  air  containing  about  2 1  volumes  of  oxygen.    In 
the  combined  state  it  forms  eight  ninths  of  water  and  nearly   • 
one  half  of  the  rocks  composing  the  earth's  crust.    It  is  also  / 
an  important  constituent  of  the  compounds  which  compose  ; 
plant  and  animal  tissues  ;  for  example,  about  66%  by  weight 
of  the  hi.  is  oxygen. 

Prepara  \!  though  oxygen  occurs  in  the  free  state 

in  the  ?  K,  its   separation  from   the  nitrogen  and 

other  u  hich  it  is  mixed  is  such  a  difficult  matter 

that  in  ;;,tory  it  has  been  found  more  convenient 

to  prept  I  its  compounds.    The  most  important  of 

the  labc  r  methods  are  the  following : 


14       AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

1 .  Preparation  from  water.  Water  is  a  compound,  consist- 
ing of   1 1. 1 8%  hydrogen  and  88.82%  oxygen.    It  is  easily 
separated  into  these  constituents  by  passing  an  electric  cur- 
rent through  it  under  suitable  conditions.    The  process  will 
be  described  in  the  chapter  on  water.    While  this  method 
of  preparation  is  a  simple  one,  it  is  not  economical. 

2.  Preparation  from  mercuric  oxide.    This  method  is  of 
interest,  since  it  is  the  one  which  led  to  the  discovery  of 
oxygen.    The  oxide,  which  consists  of   7.4%    oxygen  and 
92.6%  mercury,  is  placed  in  a  small,  glass  test  tube  and 
heated.    The   compound   is   in  this  way  decomposed  into 
mercury  which  collects  on  the  sides  of  the  glass  tube,  form- 
ing a  silvery  mirror,  and  oxygen  which,  being  a  gas,  escapes 
from  the  tube.    The  presence  of  the  oxygen  is  shown  by 
lighting  the  end  of  a  splint,  extinguishing  the  flame  and 
bringing  the  glowing  coal  into  the  mouth  of  the  tube.    The 
oxygen  causes  the  glowing  coal  to  burst  into  a  flame. 

In  a  similar  way  oxygen  may  be  obtained  from  its  compounds  with 
some  of  the  other  elements.  Thus  manganese  dioxide,  a  black  com- 
pound of  manganese  and  oxygen,  when  heated  to  about  700°,  loses  one 
third  of  its  oxygen,  while  barium  dioxide,  when  heated,  loses  one  half 
of  its  oxygen. 

3.  Preparation  from  potassium  chlorate   (usual  labora- 
tory method).    Potassium  chlorate  is  a  white   solid  which 
consists  of  31.9%  potassium,  28.9%  chlorine,  and  39.2% 
oxygen.    When  heated  it  undergoes  a  series  of  changes  in 
which  all  the  oxygen  is  finally  set  free,  leaving  a  compound 
of  potassium  and  chlorine  called  potassium  chloride.    The 
change  may  be  represented  as  follows  : 

f  potassium  "1  6  ^  1  i  ~ 

|  *        f  potassium  I  •§  -3 
^  chlorine      KI  i  =  "1    1.1     •          >%  '§  +  oxygen 

|  ||        [chlorine      i  || 
L  oxygen      J  a  °  J  3  ° 


OXYGEN 


The  evolution  of  the  oxygen  begins  at  about  400°.    It 
has  been  found,  however,  that  if  the  potassium  chlorate  is 
mixed  with  about  one  fourth  its  weight  of  manganese 
dioxide,  the  oxygen  is  given  off  at  a  much  lower  tem- 
perature.   Just  how  the  manganese  dioxide  brings 
about    this    result    is    not    definitely    known.    The 
amount  of  oxygen  obtained  from  a  given  weight  of 
potassium  chlorate  is  exactly  the  same 
whether  the  manganese  dioxide  is 
present  or  not.     So  far  as  can  be  de- 
tected the  manganese  dioxide  undergoes 

no  change. 


FIG.  4 


Directions  for 
preparing  oxy- 
gen. The  man- 
ner of  preparing 
oxygen  from  po- 
tassium chlorate 

is  illustrated  in  the  accompanying  diagram  (Fig.  4).  A  mixture  con- 
sisting of  one  part  of  manganese  dioxide  and  four  parts  of  potassium 
chlorate  is  placed  in  the  flask  A  and  gently  heated.  The  oxygen  is 
evolved  and  escapes  through  the  tube  B.  It  is  collected  by  bringing 
over  the  end  of  the  tube  the  mouth  of  a  bottle  completely  filled  with 
water  and  inverted  in  a  vessel  of  water,  as 
shown  in  the  figure.  The  gas  rises  in  the  = 
bottle  and  displaces  the  water.  In  the  prepa- 
ration of  large  quantities  of  oxygen,  a  copper 
retort  (Fig.  5)  is  often  substituted  for  the 
glass  flask. 

In  the  preparation  of  oxygen  from  potas- 
sium chlorate  and  manganese  dioxide,  the 
materials  used  must  be  pure,  otherwise  a 


FIG.  5 


violent  explosion  may  occur.    The  purity  of  the  materials  is  tested  by 
heating  a  small  amount  of  the  mixture  in  a  test  tube. 

The  collection  of  gases.    The  method  used  for  collecting  oxygen 
illustrates  the  general  method  used  for  collecting  such  gases  as  are 


1 6       AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


insoluble  in  water  or  nearly  so.  The  vessel  C  (Fig.  4),  containing 
the  water  in  which  the  bottles  are  inverted,  is  called  a  pneumatic 
trough. 

Commercial  methods  of  preparation.  Oxygen  can  now  be  pur- 
chased stored  under  great  pressure  in  strong  steel  cylinders  (Fig.  6\  It 
is  prepared  either  by  heating  a  mixture  of  potassijimchlo- 
rate  and  manganese  dioxide,  or  by  separating  it  from 
the  nitrogen  and  other  gases  with  which  it  is  mixed 
in  the  atmosphere.  The  methods  employed  for  effecting 
this  separation  will  be  described  in  subsequent  chapters. 

/ 

Physical  properties.    Oxygen  is  a  colorless, 

odorless,  tasteless  gas,  slightly  heavier  than  air. 
One  liter  of  it,  measured  at  a  temperature  of  o° 
and  under  a  pressure  of  one  atmosphere,  weighs 
1.4285  g.,  while  under  similar  conditions  one 
liter  of  air  weighs  1.2923  g.  It  is  but  slightly 
soluble  in  water.  Oxygen,  like  other  gases, 
may  be  liquefied  by  applying  very  great  pres- 
sure to  the  highly  cooled  gas.  When  the 
pressure  is  removed  the  liquid  oxygen  passes 
again  into  the  gaseous  state,  since  its  boiling 
point  under  ordinary  atmospheric  pressure  is  —  182.5°. 

Chemical  properties.  At  ordinary  temperatures  oxygen  is 
not  very  active  chemically.  Most  substances  are  either  not 
at  all  affected  by  it,  or  the  action  is  so  slow  as  to  escape 
notice.  At  higher  temperatures,  however,  it  is  very  active, 
and  unites  directly  with  most  of  the  elements.  This  activity 
may  be  shown  by  heating  various  substances  until  just  ignited 
and  then  bringing  them  into  vessels  of  the  gas,  when  they 
will  burn  with  great  brilliancy.  Thus  a  glowing  splint  intro- 
duced into  a  jar  of  oxygen  bursts  into  flame.  Sulphur 
burns  in  the  air  with  a  very  weak  flame  and  feeble  light; 
in  oxygen,  however,  the  flame  is  increased  in  size  and 


FIG.  6 


OXYGEN  17 

brightness.  Substances  which  readily  burn  in  air,  such  as 
phosphorus,  burn  in  oxygen  with  dazzling  brilliancy.  Even 
substances  which  burn  in  air  with  great  difficulty,  such  as 
iron,  readily  burn  in  oxygen. 

The  burning  of  a  substance  in  oxygen  is  due  to  the  rapid 
combination  of  the  substance  or  of  the  elements  composing 
it  with  the  oxygen.  Thus,  when  sulphur  burns  both  the 
oxygen  and  sulphur  disappear  as  such  and  there  is  formed  a 
compound  of  the  two,  which  is  an  invisible  gas,  having  the 
characteristic  odor  of  burning  sulphur.  Similarly,  phospho- 
rus on  burning  forms  a  white  solid  compound  of  phosphorus 
and  oxygen,  while  iron  forms  a  reddish-black  compound  of 
iron  and  oxygen. 

Oxidation.  The  term  oxidation  is  applied  to  the  chem- 
ical change  which  takes  place  when  a  substance,  or  one  of  its 
constituent  parts,  combines  with  oxygen.  This  process  may 
take  place  rapidly,  as  in  the  burning  of  phosphorus,  or  slowly, 
as  in  the  oxidation  (or  rusting)  of  iron  when  exposed  to  the 
air.  It  is  always  accompanied  by  the  liberation  of  heat. 
The  amount  of  heat  liberated  by  the  oxidation  of  a  definite 
weight  of  any  given  substance  is  always  the  same,  being 
entirely  independent  of  the  rapidity  of  the  process.  If  the 
oxidation  takes  place  slowly,  the  heat  is  generated  so  slowly 
that  it  is  difficult  to  detect  it.  If  the  oxidation  takes  place 
rapidly,  however,  the  heat  is  generated  in  such  a  short  inter- 
val of  time  that  the  substance  may  become  white  hot  or 
burst  into  a  flame. 

Combustion;  kindling  temperature.  When  oxidation  takes 
place  so  rapidly  that  the  heat  generated  is  sufficient  to 
cause  the  substance  to  glow  or  burst  into  a  flame  the 
process  is  called  combustion.  In  order  that  any  substance 
may  undergo  combustion,  it  is  necessary  that  it  should  be 


1 8       AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

heated  to  a  certain  temperature,  known  as  the  kindling  tem- 
perature. This  temperature  varies  widely  for  different  bodies, 
but  is  always  definite  for  the  same  body.  Thus  the  kindling 
temperature  of  phosphorus  is  far  lower  than  that  of  iron, 
but  is  definite  for  each.  When  any  portion  of  a  substance 
is  heated  until  it  begins  to  burn  the  combustion  will  con- 
tinue without  the  further  application  of  heat,  provided  the 
heat  generated  by  the  process  is  sufficient  to  bring  other 
parts  of  the  substance  to  the  kindling  temperature.  On  the 
other  hand,  if  the  heat  generated  is  not  sufficient  to  main- 
tain/the  kindling  temperature,  combustion  ceases. 

/Oxides.  The  compounds  formed  by  the  oxidation  of  any 
element  are  called  oxides.  Thus  in  the  combustion  of  sul- 
phur, phosphorus,  and  iron,  the  compounds  formed  are  called 
respectively  oxide  of  sulphur,  oxide  of  phosphorus,  and  oxide 
of  iron.  In  general,  then,  an  oxide  is  a  compound  of  oxygen 
ivith  another  clement.  A  great  many  substances  of  this 
class  are  known ;  in  fact,  the  oxides  of  all  the  common  ele- 
ments have  been  prepared,  with*  the  exception  of  those  of 
fluorine  and  bromine.  Some  of  these  are  familiar  com- 
pounds. Water,  for  example,  is  an  oxide  of  hydrogen,  and 
lime  an  oxide  of  the  metal  calcium. 

Products  of  combustion.    The  particular  oxides   formed 
by  the  combustion  of  any  substance  are  called  products  of 
combustion  of  that  substance.    Thus  oxide  of  sulphur  is  the 
product  of  the  combustion  of  sulphur ;  oxide  of  iron  is  the 
product   of   the   combustion    of   iron.     It    is    evident   that 
the  products  of  the  combustion  of  any  substance  must  weigh 
more  than  the  original  substance,  the  increase  in  weight 
corresponding  to  the  amount  of  oxygen  taken  up  in  the  act  1 
of  combustion.    For  example,  when  iron  burns  the  oxide  of  I 
iron  formed  weighs  more  than  the  original  iron. 


OXYGEN 


In  some  cases  the  products  of  combustion  are  invisible 
gases,  so  that  the  substance  undergoing  combustion  is 
apparently  destroyed.  Thus,  when  a  candle  burns  it  is  con- 
sumed, and  so  far  as  the  eye  can  judge  nothing  is  formed 
during  combustion.  That  invisible  gases  are  formed,  how- 
ever, and  that  the  weight  of  these  is  greater  than  the  weight 
of  the  candle  may  be  shown  by  the  following  experiment. 

A  lamp  chimney  is  filled  with  sticks  of  the  compound  known  as 
sodium  hydroxide  (caustic  soda),  and  suspended  from  the  beam  of  the 
balance,  as  shown  in  Fig.  7.  A  piece 
of  candle  is  placed  on  the  balance 
pan  so  that  the  wick  comes  just  below 
the  chimney,  and  the  balance  is 
brought  to  a  level  by  adding  weights 
to  the  other  pan.  The  candle  is  then 
lighted.  The  products  formed  pass 
up  through  the  chimney  and  are 
absorbed  by  the  sodium  hydroxide. 
Although  the  candle  burns  away,  the 
pan  upon  which  it  rests  slowly  sinks, 
showing  that  the  combustion  is  at- 
tended by  an  increase  in  weight. 


FIG.  7 


Combustion  in  air  and  in  oxygen.  Combustion  in  air  and  in  oxygen 
differs  only  in  rapidity,  the  products  formed  being  exactly  the  same. 
That  the  process  should  take  place  less  rapidly  in  the  former  is  readily 
understood,  for  the  air  is  only  about  one  fifth  oxygen,  the  remaining 
four  fifths  being  inert  gases.  Not  only  is  less  cxygen  available,  but 
much  of  the  heat  is  absorbed  in  raising  the  temperature  of  the  inert 
gases  surrounding  the  substance  undergoing  combustion,  and  the  tem- 
perature reached  in  the  combustion  is  therefore  less. 

Phlogiston  theory  of  combustion.^  The  French  chemist  Lavoisier 
(1743-1794),  who  gave  to  oxygen  its  name,  was  the  first  to  show  that 
combustion  is  due  to  union  with  oxygen.  Previous  to  his  time  com- 
bustion was  supposed  to  be  due  to  the  presence  of  a  substance  or 
principle  called  phlogiston.  One  substance  was  thought  to  be  more 
combustible  than  another  because  it  contained  more  phlogiston.  Coal, 
for  example,  was  thought  to  be  very  rich  in  phlogiston.  The  ashes 


20       AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

left  after  combustion  would  not  burn  because  all  the  phlogiston  had 
escaped.  If  the  phlogiston  could  be  restored  in  any  way,  the  substance 
would  then  become  combustible  again.  Although  this  view  seems 
absurd  to  us  in  the  light  of  our  present  knowledge,  it  formerly  had 
general  acceptance.  The  discovery  of  oxygen  led  Lavoisier  to  investi- 
gate the  subject,  and  through  his  experiments  he  arrived  at  the  true 
explanation  of  combustion.  The  discovery  of.  oxygen  together  with 
the  part  it  plays  in  combustion  is  generally  regarded  as  the  most 
important  discovery  in  the  history  of  chemistry.  It  marked  the  dawn 
of  a  new  period  in  the  growth  of  the  science. 

Combustion  in  the  broad  sense.  According  to  the  defini- 
tion given  above,  the  presence  of  oxygen  is  necessary  for 
combustion.  The  term  is  sometimes  used,  however,  in  a 
broader  sense  to  designate  any  chemical  change  attended 
by  the  evolution  of  heat  and  light.  Thus  iron  and  sulphur, 
or  hydrogen  and  chlorine  under  certain  conditions,  will  com- 
bine so  rapidly  that  light  is  evolved,  and  the  action  is  called 
a  combustion.  Whenever  combustion  takes  place  in  the  air, 
however,  the  process  is  one  of  oxidation. 

Spontaneous  combustion.  The  temperature  reached  in  a  given  chemi- 
cal action,  such  as  oxidation,  depends  upon  the  rate  at  which  the 
reaction  takes  place.  This  rate  is  usually  increased  by  raising  the 
temperature  of  the  substances  taking  part  in  the  action. 

When  a  slow  oxidation  takes  place  under  such  conditions  that  the 
heat  generated  is  not  lost  by  being  conducted  away,  the  temperature 
of  the  substance  undergoing  oxidation  is  raised,  and  this  in  turn 
hastens  the  rate  of  oxidation.  The  rise  in  temperature  may  continue 
in  this  way  until  the  kindling  temperature  of  the  substance  is  reached, 
when  combustion  begins.  Combustion  occurring  in  this  way  is  called 
spontaneous  combustion. 

Certain  oils,  such  as  the  linseed  oil  used  in  paints,  slowly  undergo 
oxidation  at  ordinary  temperatures,  and  not  infrequently  the  origin  of 
fires  has  been  traced  to  the  spontaneous  combustion  of  oily  rags.  The 
spontaneous  combustion  of  hay  has  been  known  to  set  barns  on  fire. 
Heaps  of  coal  have  been  found  to  be  on  fire  when  spontaneous  com- 
bustion offered  the  only  possible  explanation. 


OXYGEN  2 1 

Importance  of  oxygen,  i.  Oxygen  is  essential  to  life. 
Among  living  organisms  only  certain  minute  forms  of  plant 
life  can  exist  without  it.  In  the  process  of  respiration  the 
air  is  taken  into  the  lungs  where  a  certain  amount  of  oxygen 
is  absorbed  by  the  blood.  It  is  then  carried  to  all  parts  of  the 
body,  oxidizing  the  worn-out  tissues  and  changing  them  into 
substances  which  may  readily  be  eliminated  from  the  body. 
The  heat  generated  by  this  oxidation  is  the  source  of 
the  heat  of  the  body.  The  small  amount  of  oxygen  which 
water  dissolves  from  the  air  supports  all  the  varied  forms 
of  aquatic  animals. 

2.  Oxygen  is  also  essential  to  decay.  The  process  of  decay 
is  really  a  kind  of  oxidation,  but  it  will  only  take  place  in 
the  presence  of  certain  minute  forms  of  life  known  as  bac- 
teria.   Just  how  these  assist  in  the  oxidation  is  not  known. 
By  this  process  the  dead  products  of  animal  and  vegetable 
life  which  collect  on  the  surface  of  the  earth  are  slowly  oxi- 
dized and  so  converted  into  harmless  substances.    In  this 
way  oxygen  acts  as  a  great  purifying  agent. 

3.  Oxygen  is  also  used  in  the  treatment  of  certain  dis- 
eases in  which  the  patient  is  unable  to  inhale  sufficient  air 
to  supply  the  necessary  amount  of  oxygen. 


OZONE 

Preparation.  When  electric  sparks  are  passed  through  oxygen  or 
air  a  small  percentage  of  the  oxygen  is  converted  into  a  substance 
called  ozone,  which  differs  greatly  from  oxygen  in  its  properties.  The 
same  change  can  also  be  brought  about  by  certain  chemical  processes. 
Thus,  if  some  pieces  of  phosphorus  are  placed  in  a  bottle  and  partially 
covered  with  water,  the  presence  of  ozone  may  soon  be  detected  in  the 
air  contained  in  the  bottle.  The  conversion  of  oxygen  into  ozone  is 
attended  by  a  change  in  volume,  3  volumes  of  oxygen  forming  2  vol- 
umes of  ozone.  If  the  resulting  ozone  is*  heated  to  about  300°,  the 


22        AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

reverse  change  takes  place,  the  2  volumes  of  ozone  being  changed 
back  into  3  volumes  of  oxygen.  It  is  possible  that  traces  of  ozone 
exist  in  the  atmosphere,  although  its  presence  there  has  not  been 
definitely  proved,  the  tests  formerly  used  for  its  detection  having  been 
shown  to  be  unreliable. 

Properties.  As  commonly  prepared,  ozone  is  mixed  with  a  large 
excess  of  oxygen.  It  is  possible,  however,  to  separate  the  ozone  and 
thus  obtain  it  in  pure  form.  The  gas  so  obtained  has  the  character- 
istic odor  noticed  about  electrical  machines  when  in  operation.  By 
subjecting  it  to  great  pressure  and  a  low  temperature,  the  gas  con- 
denses to  a  bluish  liquid,  boiling  at  —  1 19°.  When  unmixed  with  other 
gases  ozone  is  very  explosive,  changing  back  into  oxygen  with  the 
liberation  of  heat.  Its  chemical  properties  are  similar  to  those  of 
oxygen  except  that  it  is  far  more  active.  Air  or  oxygen  containing 
a  small  amount  of  ozone  is  now  used  in  place  of  oxygen  in  certain 

\\  ^manufacturing  processes. 
)r       The  difference  between  oxygen  and  ozone.    Experiments  show  that 

'/  ^  in  changing  oxygen  into  ozone  no  other  kind  of  matter  is  either  added 
to  the  oxygen  or  withdrawn  from  it.  The  question  arises  then,  How  can 
we  account  for  the  difference  in  their  properties  ?  It  must  be  remem- 
bered that  in  all  changes  we  have  to  take  into  account  energy  as  well 
as  matter.  By  changing  the  amount  of  energy  in  a  substance  we 
change  its  properties.  That  oxygen  and  ozone  contain  different 
amounts  of  energy  may  be  shown  in  a  number  of  ways ;  for  example, 
by  the  fact  that  the  conversion  of  ozone  into  oxygen  is  attended  by  the 
liberation  of  heat.  The  passage  of  the  electric  sparks  through  oxygen 
has  in  some  way  changed  the  energy  content  of  the  element  and  thus 
it  has  acquired  new  properties.  Oxygen  and  ozone  must,  therefore,  be 
regarded  as  identical  so  far  as  the  kind  of  matter  of  which  they  are 
composed  is  concerned.  Their  different  properties  are  due  to  their  dif- 
ferent energy  contents. 

Allotropic  states  or  forms  of  matter.  Other  elements  besides  oxygen 
may  exist  in  more  than  one  form.  These 'different  forms  of  the  same 
element  are  called  allotropic  states  or  forms  of  the  element.  These 
forms  differ  not  only  in  physical  properties  but  also  in  their  energy 
contents.  Elements  often  exist  in  a  variety  of  forms  which  look  quite 
different.  These  differences  may  be  due  to  accidental  causes,  such  as 
the  size  or  shape  of  the  particles  or  the  way  in  which  the  element  was 
prepared.  Only  such  forms,  however,  as  have  different  energy  contents 
are  properly  called  allotropic  forms. 


OXYGEN  23 

MEASUREMENT  OF  GAS  VOLUMES 

Standard  conditions.  It  is  a  well-known  fact  that  the  volume  occu- 
pied by  a  definite  weight  of  any  gas  can  be  altered  by  changing  the 
temperature  of  the  gas  or  the  pressure  to  which  it  is  subjected.  In 
measuring  the  volume  of  gases  it  is  therefore  necessary,  for  the  sake 
of  accuracy,  to  adopt  some  standard  conditions  of  temperature  and 
pressure.  The  conditions  agreed  upon  are  (i)  a  temperature  of  o°, 
and  (2)  a  pressure  equal  to  the  average  pressure  exerted  by  the  at- 
mosphere at  the  sea  level,  that  is,  1033.3  g.  per  square  centimeter. 
These  conditions  of  temperature  and  pressure  are  known  as  the 
standard  conditions,  and  when  the  volume  of  a  gas  is  given  it  is  under- 
stood that  the  measurement  was  made  under  these  conditions,  unless 
it  is  expressly  stated  otherwise.  For  example,  the  weight  of  a  liter 
of  oxygen  has  been  given  as  1.4285  g.  This  means  that  one  liter 
of  oxygen,  measured  at  a  temperature  of  o°  and  under  a  pressure  of 
1033. 3  g.  per  square  centimeter,  weighs  1.4285  g. 

The  conditions  which  prevail  in  the  laboratory  are  never  the  stand- 
ard conditions.  It  becomes  necessary,  therefore,  to  find  a  way  to  cal- 
culate the  volume  which  a  gas  will  occupy  under  standard  conditions 
from  the  volume  which  it  occupies  under  any  other  conditions.  This 
may  be  done  in  accordance  with  the  following  laws. 

Law  of  Charles.  This  law  expresses  the  effect  which  a  change  in 
the  temperature  of  a  gas  has  upon  its  volume.  It  may  be  stated  as 
follows :  For  every  degree  the  temperature  of  a  gas  rises  above  zero 
the  volume  of  the  gas  is  increased  by  273-  of  the  volume  which  it  occu- 
pies at  zero  ;  likewise  for  every  degree  the  temperature  of  the  gas  falls . 
below  zero  the  volume  of  the  gas  is  decreased  by  -%\^  of  the  volume 
which  it  occupies  at  zero,  provided  in  both  cases  that  the  pressure  to 
which  the  gas  is  subjected  remains  constant.  "  ^— -^ 

If  /^represents  the  volume  of  gas  at  o°,  then  the  volume  aVi0  will 
be  V+  ^  V\  3t  7°  it  will  be  V+  •$-$  V;  or,  in  general,  the  volume 
V)  at  the  1  t,  will  be  expressed  by  the  formula 

v=  V  +  —  V, 
273 


Since  the  formula  may  be  written 

v-V(\  +  0.00366  /). 


24       AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Since  the  value  of  V  (volume  under  standard  conditions)  is  the  one 
usuajjj:  sought,  it  is  convenient  to  transpose  the  equation  to  the  fol- 
lowing form  : 

^      V=  i  +  0.00366 1' 

The  following  problem  will  serve  as  an  illustration  of  the  applica- 
tion of  this  equation. 

The  volume  of  a  gas  at  20°  is  750  cc. ;  find  the  volume  it  will  occupy 
at  o°,  the  pressure  remaining  constant. 

In  this  case,  v  =  750  cc.  and  /  =  20.  By  substituting  these  values, 
equation  (4)  becomes 

V— Ld —  698.9  cc. 

i  -f  0.00366  x  20 

Law  of  Boyle.  This  law  expresses  the  relation  between  the  volume 
occupied  by  a  gas  and  the  pressure  to  which  it  is  subjected.  It  may 
be  stated  as  follows :  The  volume  of  a  gas  is  inversely  proportional 
to  the  pressure  under  which  it  is  measured,  provided  the  temperature 
of  the  gas  remains  constant. 

If  V  represents  the  volume  when  subjected  to  a  pressure  P  and  v 
represents  its  volume  when  the  pressure  is  changed  to  p,  then,  in 
accordance  with  the  above  law,(  V:  'v  :  :p  :  P,  or  VP  —  vp.  In  other 
words,  for  a  given  weight  of  a  gas  the  product  of  the  numbers  repre- 
senting its  volume  and  the  pressure  to  which  it  is  subjected  is  a 
constant. 

Since  the  pressure  of  the  atmosphere  at  any  point  is  indicated  by 
the  barometric  reading,  it  is  convenient  in  the  solution  of  the  problems 
to  substitute  the  latter  for  the  pressure  measured  in  grams  per  square 
centimeter.  The  average  reading  of  the  barometer  at  the  sea  level  is 
760  mm.,  which  corresponds  to  a  pressure  of  1033.3  g-  Per  square  cen- 
timeter. The  following  problem  will  serve  as  an  illustration  of  the 
application  of  Boyle's  law. 

A  gas  occupies  a  volume  of  500  cc.  in  a  laboratory  where  the  baro- 
metric reading  is  740  mm.  What  volume  would  it  occupy  if  the 
atmospheric  pressure  changed  so  that  the  reading  became  750  mm.? 

Substituting  the  values  in  the  equation  VP  =  vp,  we  have  500  X 
740  =  v  x  750,  or  v  —  493.3  cc. 

Variations  in  the  volume  of  a  gas  due  to  changes  both  in  tempera- 
ture and  pressure.  Inasmuch  as  corrections  must  be  made  as  a  rule 


OXYGEN  25 

for  both  temperature  and  pressure,  it  is  convenient  to  combine  the 
equations  given  above  for  the  corrections  for  each,  so  that  the,  two 
corrections  may  be  made  in  one  operation.  The  following  equation  is 
thus  obtained : 

V  —  ^ 

=  760(1  4-  0.00366  /)' 

in  which  Vs  represents  the  volume  of  a  gas  under  standard  conditions 
and  v,  p,  and  t  the  volume,  pressure,  and  temperature  respectively  at 
which  the  gas  was  actually  measured. 

The  following  problem  will  serve  to  illustrate  the  application  of  this 
equation. 

A  gas  having  a  temperature  of  20°  occupies  a  volume  of  500  cc. 
when  subjected  to  a  pressure  indicated  by  a  barometric  reading  of 
740  mm.  What  volume  would  this  gas  occupy  under  standard  con- 
ditions ? 

In  this  problem  v  =  500,  p  =  740,  and  /=  20.  Substituting  these 
values  in  the  above  equation,  we  get 

coo  x  740 

Vs  = 1 — =453-6  cc. 

760  (i  +  0.00366  x  20) 

Variations  in  the  volume  of  a  gas  due  to  the  pressure  of  aqueous 
vapor.  In  many  cases  gases  are  collected  over  water,  as  explained 
under  the  preparation  of  oxygen.  In  such  cases  there 
is  present  in  the  gas  a  certain  amounjt  of  water  vapor. 
This  vapor  exerts  a  definite  pressure,  which  acts  in  oppo- 
sition to  the  atmospheric  pressure  and  which  therefore 
must  be  subtracted  from  the  latter  in  determining  the 
effective  pressure  upon  the  gas.  Thus,  suppose  we  wish 
to  determine  the  pressure  to  which  the  gas  in  tube  A 
(Fig.  8)  is  subjected.  The  tube  is  raised  or  lowered 
until  the  level  of  the  water  inside  and  outside  the  tube 
is  the  same.  The  atmosphere  presses  down  upon  the 
surface  of  the  water  (as  indicated  by  the  arrows),  thus 
forcing  the  water  upward  within  the  tube  with  a  pressure 
equal  to  the  atmospheric  pressure.  The  full  force  of 
this  upward  pressure,  however,  is  not  spent  in  com- 
pressing the  gas  within  the  tube,  for  since  it  is  collected  over  water  it 
contains  a  certain  amount  of  water  vapor.  This  water  vapor  exerts  a 
pressure  (as  indicated  by  the  arrow  within  the  tube)  in  opposition  to 


26       AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

the  upward  pressure.  It  is  plain,  therefore,  that  the  effective  pressure 
upon  the  gas  is  equal  to  the  atmospheric  pressure  less  the  pressure 
exerted  by  the  aqueous  vapor.  The  pressure  exerted  by  the  aqueous 
vapor  increases  with  the  temperature.  The  figures  representing  the 
extent  of  this  pressure  (often  called  the  tension  of  aqueous  vapor)  are 
given  in  the  Appendix.  They  express  the  pressure  or  tension  in  milli- 
meters of  mercury,  just  as  the  atmospheric  pressure  is  expressed  in 
millimeters  of  mercury.  Representing  the  pressure  of  the  aqueous 
vapor  by  a,  formula  (5)  becomes 

(6)     K  =  -      *0»-*) 

760  (i  -f  0.00366  /) 

The  following  problem  will  serve  to  illustrate  the  method  of  applying 
the  correction  for  the  pressure  of  the  aqueous  vapor. 

The  volume  of  a  gas  measured  over  water  in  a  laboratory  where 
the  temperature  is  20°  and  the  barometric  reading  is  740  mm.  is  500  cc. 
What  volume  would  this  occupy  under  standard  conditions  ? 

The  pressure  exerted  by  the  aqueous  vapor  at  20°  (see  table  in 
Appendix)  is  equal  to  the  pressure  exerted  by  a  column  of  mercury 
17.4  mm.  in  height.  Substituting  the  values  of  v,  t,  p,  and  a  in  for- 
mula (6),  we  have 

V,  =        5°o(74°-l74)        =  cc 

760  (i  +  0.00366  x  20) 

Adjustment  of  tubes  before  reading  gas  volumes.  In  measuring  the 
volumes  of  gases  collected  in  graduated  tubes  or  other  receivers,  over 
a  liquid  as  illustrated  in  Fig.  8,  the  reading  should  be  taken  after 
raising  or  lowering  the  tube  containing  the  gas  until  the  level  of  the 
liquid  inside  and  outside  the  tube  is  the  same  ;  for  it  is  only  under 
these  conditions  that  the  upward  pressure  within  the  tube  is  the  same 
as  the  atmospheric  pressure. 


EXERCISES 

1.  What  is  the  meaning  of  the  following  words  ?  phlogiston,  ozone, 
phosphorus.     (Consult  dictionary.) 

2.  Can  combustion  take  place  without  the  emission  of  light? 

3.  Is  the  evolution  of  light  always  produced  by  combustion  ? 

4.  (a)  What  weight  of  oxygen  can  be  obtained  from  100  g.  of 
water  ?  (b)  What  volume  would  this  occupy  under  standard  conditions  ? 


OXYGEN  27 

v  5.  (a)  What  weight  of  oxygen  can  be  obtained  from  500  g.  of 
mercuric  oxide  ?  (b)  What  volume  would  this  occupy  under  standard 
conditions? 

6.  What  weight  of  each  of  the  following  compounds  is  necessary 
to  prepare  50  1.  of  oxygen  ?  (a)  water  ;  (^)  mercuric  oxide  ;  (c}  potas- 
sium chlorate. 

7.  Reduce  the  following  volumes  to  o°,  the  pressure  remaining 
constant:  (a)  150  cc.  at  10°;  (£)  840  cc.  at  273°. 

I.  A  certain  volume  of  gas  is  measured  when  the  temperature  is 
20°.  At  what  temperature  will  its  volume  be  doubled  ? 

t.  Reduce  the  following  volumes  to  standard  conditions  of  pres- 
sure, the  temperature  remaining  constant :  (a)  200  cc.  at  740  mm. ; 
(£)  500  1.  at  380  mm. 

It.    What  is  the  weight  of    i  1.  of  oxygen  when  the  pressure  is 
750  mm.  and  the  temperature  o°? 

11.  Reduce  the  following  volumes  to  standard  conditions  of  tem- 
perature and  pressure :  (a)  340  cc.  at  12°  and  753  mm  ;  (£)  500  cc.  at 
15°  and  740  mm. 

12.  What  weight  of   potassium  chlorate  is  necessary  to  prepare 
250  1.  of  oxygen  at  20°  and  750  mm.  ? 

13.  Assuming  the  cost  of  potassium  chlorate  and  mercuric  oxide 
to  be  respectively  $0.50  and  $1.50  per  kilogram,  calculate  the  cost  of 
materials  necessary  for  the  preparation  of  50  1.  of  oxygen  from  each  of 
the  above  compounds. 

14.  100  g.  of  potassium  chlorate  and  25  g.  of  manganese  dioxide 
were  heated  in  the  preparation  of  oxygen.    What  products  were  left 
in  the  flask,  and  how  much  of  each  was  present? 


CHAPTER  III 

i 

HYDROGEN 

i 

Historical.  The  element  hydrogen  was  first  clearly  rec- 
ognized as  a  distinct  substance  by  the  English  investiga- 
tor Cavendish,  who  in  1 766  obtained  it  in  a  pure  state,  and 
showed  it  to  be  different  from  the  other  inflammable  aip  or 
gases  which  had  long  been  known.  Lavoisier  gave  if  the 
name  hydrogen,  signifying  water  former,  since  it  had  been 
found  to  be  a  constituent  of  water. 

Occurrence.  In  the  free  state  hydrogen  is  found  in  the 
atmosphere,  but  only  in  traces.  In  the  combined  state  it  is 
widely  distributed,  being  a  constituent  of  water  as  well  as 
of  all  living  organisms,  and  the  products  derived  from  them, 
such  as  starch  and  sugar.  About  10%  of  the  human  body 
is  hydrogen.  Combined  with  carbon,  it  forms  the  sub- 
stances which  constitute  petroleum  and  natural  gas. 

It  is  an  interesting  fact  that  while  hydrogen  in  the  free  state  occurs 
only  in  traces  on  the  earth,  it  occurs  in  enormous  quantities  in  the 
gaseous  matter  surrounding  the  sun  and  certain  other  stars. 

Preparation  from  water.  Hydrogen  can  be  prepared  from 
water  by  several  methods,  the  most  important  of  which  are 
the  following. 

1.  By  the  electric  current.   .As  has  been  indicated  in  the 
preparation  of  oxygen,  water  is  easily  separated  into  its 
constituents,   hydrogen  and  oxygen,  by  passing  an  electric 
current  through  it  under  certain  conditions. 

2.  By  the  action  of  certain  metals!  When  brought  into 
contact  with  certain  metals  under  appropriate  conditions, 

28 


HYDROGEN  29 

water  gives  up  a  portion  or  the  whole  of  its  hydrogen,  its  place 
being  taken  by  the  metal.  In  the  case  of  a  few  of  the  metals 
this  change  occurs  at  ordinary  temperatures.  Thus,  if  a  bit 
of  sodium  is  thrown  on  water,  an  action  is  seen  to  take  place 
at  once,  sufficient  heat  being  generated  to  melt  the  sodium, 
which  runs  about  on  the  surface  of  the  water.  The  change 
which  takes  place  consists  in  the  displacement  of  one  half 
of  the  hydrogen  of  the  water  by  the  sodium,  and  may  be 
represented  as  follows : 

f  hydrogeji .  "I  f  sodium      "]  g  4f 

sodium  +  •<  hydrogen  f  |  *=  «|  hydrogen  l|  |  +  hydrogen 
[  oxygen  « .  J  >        [  oxygen      J  ~I 

The  sodium  hydroxide  formed  is  a  white  solid  which 
remains  dissolved  in  the  undecomposed  water,  and  may  be 
obtained  by  evaporating  the  solution  to  dryness.  The  hydro- 
gen is  evolved  as  a  gas  and  may  be  collected  by  suitable 
apparatus. 

Other  metals,  such  as  magnesium  and  iron,  decompose 
water  rapidly,  but  only  at  higher  temperatures.  When 
steam  is  passed  over  hot  iron,  for  example,  the  iron  com- 
bines with  the  oxygen  of  the  steam,  thus  displacing  the 
hydrogen.  Experiments  show  that  the  change  may  be  rep- 
resented as  follows : 

t 

f  hydrogen]  _  , 

£       f  iron       I  a  -&      f  hydrogen  j 
iron  +  4  hydrogen  f  -s  =  S  r  J  S  +  S  u  _j 

[  I      1  oxygen  J  -  g      \  hydrogen  J 
[oxygen      J 

The  iron  oxide  formed  is  a  reddish-black  compound,  identi- 
cal with  that  obtained  by  the  combustion  of  iron  in  oxygen. 

Directions  for  preparing  hydrogen  by  the  action  of  steam  on  iron. 

The  apparatus  used  in  the  preparation  of  hydrogen  from  iron  and 


AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


steam  is  shown  in  Fig.  9.  A  porcelain  or  iron  tube  B,  about  50  cm.  in 
length  and  2  cm.  or  3  cm.  in  diameter,  is  partially  filled  with  fine  iron 
wire  or  tacks  and  connected  as  shown  in  the  figure.  The  tube  B  is 
heated,  slowly  at  first,  until  the  iron  is  red-hot.  Steam  is  then  con- 
ducted through  the  tube  by  boiling  the  water  in  the  flask  A.  The  hot 
iron  combines  with  the  oxygen  in  the  steam,  setting  free  the  hydrogen, 
which  is  collected  over  water.  The  gas  which  first  passes  over  is 
mixed  with  the  air  previously  contained  in  the  flask  and  tube,  and  is 
allowed  to  escape,  since  a  mixture  of  hydrogen  with  oxy- 
gen or  air  explodes  violently  when  brought  in  contact  with 
a  flctme.  It  is  evident  that  the  flask  A  must  be  discon- 
nected from  the 
tube  before  the 
heat  is  with- 
drawn. 

That  the  gas 
obtained  is  dif- 


FIG.  9 

ferent  from  air  and  oxygen  may  be  shown  by  holding  a  bottle  of  it 
mouth  downward  and  bringing  a  lighted  splint  into  it.  The  hydrogen 
is  ignited  and  burns  with  an  almost  colorless  flame. 

Preparation  from  acids  (tisual  laboratory  method}.  While 
hydrogen  can  be  prepared  from  water,  either  by  the  action 
of  the  electric  current  or  by  the  action  of  certain  metals, 
these  methods  are  not  economical  and  are  therefore  but 
little  used.  In  the  laboratory  hydrogen  is  generally  pre- 
pared from  compounds  known  as  acids,  all  of  which  con- 
tain hydrogen.  When  acids  are  brought  in  contact  with 
certain  metals,  the  metals  dissolve  and  set  free  the  hydrogen 


HYDROGEN  3 1 

of  the  acid.  Although  this  reaction  is  a  quite  general  one, 
it  has  been  found  most  convenient  in  preparing  hydrogen 
by  this  method  to  use  either  zinc  or  iron  as  the  metal  and 
either  hydrochloric  or  sulphuric  acid  as  the  acid.  Hydro- 
chloric acid  is  a  compound  consisting  of  2.77%  hydrogen 
and  97.23%  chlorine,  while  sulphuric  acid  consists  of  2.05% 
hydrogen,  32.70%  sulphur,  and  65.25%  oxygen. 

The  changes  which  take  place  in  the  preparation  of 
hydrogen  from  zinc  and  sulphuric  acid  (diluted  with  water) 
may  be  represented  as  follows  : 

f  hydrogen  1  ,a  f  zinc        1      ^ 

zinc  +  J  sulphur      L|,|  =  J  sulphur  XjJ  +  hydrogen 
[oxygen      j  $,          [oxygen  J      *    . 

In  other  words,  the  zinc  has  taken  the  place  of  the  hydro- 
gen in  sulphuric  acid.  The  resulting  compound  contains 
zinc,  sulphur,  and  oxygen,  and  is  known  as  zinc  sulphate. 
This  remains  dissolved  in  the  water  present  in  the  acid.  It 
may  be  obtained  in  the  form  of  a  white  solid  by  evaporating 
the  liquid  left  after  the  metal  has  passed  into  solution. 

When  zinc  and  hydrochloric  acid  are  used  the  following 
changes  take  place : 

f  hydrogen!  £  «  ^       fzinc        1  o  3 
zinc  +  4    :,-    .         r^Sl^^    ui     •       rl  °  -h  hydrogen 
[chlorine    J  e2  3        [  chlorine  J  ~2 

When  iron  is  used  the  changes  which  take  place  are 
exactly  similar  to  those  just  given  for  zinc. 

Directions  for  preparing  hydrogen  from  acids.  The  preparation  of 
hydrogen  from  acids  is  carried  out  in-the  laboratory  as  follows:  The 
metal  is  placed  in  a  flask  or  wide-mouthed  bottle  A  (Fig.  10)  and  the 
acid  is  added  slowly  through  the  funnel  tube  B.  The  metal  dissolves 
in  the  acid,  while  the  hydrogen  which  is  liberated  escapes  through  the 
exit  tube  Cand  is  collected  over  water.  It  is  evident  that  the  hydrogen 


32        AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


0 


which  passes  over  first  is  mixed  with  the  air  from  the  bottle  A. 
Hence  care  must  be  taken  not  to  bring  a  flame  near  the  exit  tube, 
since,  as  has  been  stated  previously,  such  a  mixture  explodes  with 

great  violence  when  brought  in  contact 
with  a  flame. 

Precautions.  Both  sulphuric  acid  and 
zinc,  if  impure,  are  likely  to  contain 
small  amounts  of  arsenic.  Such  mate- 
rials should  not  be  used  in  preparing 
hydrogen,  since  the  arsenic  present  com- 
bines with  a  portion  of  the  hydrogen  to 


B 


FIG. 10 

form  a  very  poisonous  gas  known  as  arsine.  On  the  other  hand, 
chemically  pure  sulphuric  acid,  i.e.  sulphuric  acid  that  is  entirely  free 
from  impurities,  will  not  act  upon  chemically  pure  zinc.  The  reaction 
may  be  started,  however,  by  the  addition  of  a  few  drops  of  a  solution 
of  copper  sulphate  or  platinum  tetrachloride. 

Physical  properties.  Hydrogen  is  similar  to  oxygen  in 
that  it  is  a  colorless,  tasteless,  odorless  gas.  It  is  character- 
ized by  its  extreme  lightness,  being  the  lightest  of  all  known 
substances.  One  liter  of  the  gas  weighs  only  0.08984  g. 
On  comparing  this  weight  with  that  of  an  equal  volume 
of  oxygen,  viz.,  1.4285  g.,  the  latter  is  found  to  be  15.88 
times  as  heavy  as  hydrogen.  Similarly,  air  is  found  to  be 
14.38  times  as  heavy  as  hydrogen.  Soap  bubbles  blown 
with  hydrogen  rapidly  rise  in  the  air.  On  account  of  its 
lightness  it  is  possible  to  pour  it  upward  from  one  bottle 
into  another.  Thus,  if  the  bottle  A  (Fig.  n)  is  filled  with 
hydrogen,  placed  mouth  downward  by  the  side  of  bottle  B, 


HYDROGEN 


33 


filled  with  air,  and  is  then  gradually  inverted  under  B  as 
indicated  in  the  figure,  the  hydrogen  will  flow  upward  into 
bottle  B,  displacing  the  air.  Its  presence  in  bottle  B  may 
then  be  shown  by  bringing  a  lighted  splint  to  the  mouth  of 
the  bottle,  when  the  hydrogen  will  be 
ignited  by  the  flame.  It  is  evident,  from 
this  experiment,  that  in  order  to  retain 
the  gas  in  an  open  bottle  the  bottle 
must  be  placed  mouth  downward. 

Hydrogen    is   far   more    difficult  to 
liquefy  than  any  other  gas,  with  the  FlG-  1  l 

exception  of  the  rare  element  helium,  which  has  as  yet 
resisted  all  efforts  to  liquefy  it.  The  English  scientist 
Dewar,  however,  in  1898  succeeded  not  only  in  obtaining 
hydrogen  in  liquid  state  but  also  as  a  solid.  Liquid  hydro- 
gen is  colorless  and  has  a  density  of  only  0.07.  Its  boiling 
point  under  atmospheric  pressure  is  —  252°.  Under  dimin- 
ished pressure  the  temperature  has  been  reduced  to—  262°. 
The  solubility  of  hydrogen  in  water  is  very  slight,  being 
still  less  than  that  of  oxygen. 

Pure  hydrogen  produces  no  injurious  results  when  inhaled. 
Of  course  one  could  not  live  in  an  atmosphere  of  the  gas, 
since  oxygen  is  essential  to  respiration. 

Chemical  properties.  At  ordinary  temperatures  hydro- 
gen is  not  an  active  element.  A  mixture  of  hydrogen  and 
chlorine,  however,  will  combine  with  explosive  violence  at 
ordinary  temperature  if  exposed  to  the  sunlight.  The  union 
can  be  brought  about  also  by  heating.  The  product  formed 
in  either  case  is  hydrochloric  acid.  Under  suitable  condi- 
tions hydrogen  combines  with  nitrogen  to  form  ammonia,  and 
with  sulphur  to  form  the  foul-smelling  •  gas,  hydrogen  sul- 
phide. The  affinity  of  hydrogen  for  oxygen  is  so  great  that 


34 


AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


a  mixture  of  hydrogen  and  oxygen  or  hydrogen  and  air 
explodes  with  great  violence  when  heated  to  the  kindling 
temperature  (about  612°),  Nevertheless  under  proper  con- 
ditions hydrogen  may  be  made  to  burn  quietly  in  either 
oxygen  or  air.  The  resulting  hydrogen  flame  is  almost 
colorless  and  is  very  hot.  The  combustion  of  the  hydrogen 
is,  of  course,  due  to  its  union  with  oxygen.  The  product 
of  the  combustion  is  therefore  a  compound  of  hydrogen 

and  oxygen.  That 
this  compound  is 
water  may  be 
shown  easily  by 
experiment. 

Directions  for 
burning  hydrogen  in 
air.  The  combustion 
of  hydrogen  in  air 
may  be  carried  out 
safely  as  follows: 
The  hydrogen  is  gen- 
erated in  the  bottle  A 
(Fig.  1 2),  is  dried  by 
conducting  it  through  the  tube  X,  filled  with  some  substance  (gener- 
ally calcium  chloride)  which  has  a  great  attraction  for  moisture,  and 
escapes  through  the  tube  T,  the  end  of  which  is  drawn  out  to  a  jet. 
The  hydrogen  first  liberated  mixes  with  the  air  contained  in  the  gen- 
erator. If  a  flame  is  brought  near  the  jet  before  this  mixture  has  all 
escaped,  a  violent  and  very  dangerous  explosion  results,  since  the 
entire  apparatus  is  filled  with  the  explosive  mixture.  On  the  other 
hand,  if  the  flame  is  not  ^applied  until  all  the  air  has  been  expelled, 
the  hydrogen  is  ignited  and  burns  quietly,  since  only  the  small  amount 
of  it  which  escapes  from  the  jet  can  come  in  contact  with  the  oxygen 
of  the  air  at  any  one  time.  By  holding  a  cold,  dry  bell  jar  or  bottle 
over  the  flame,  in  the  manner  shown  in  the  figure,  the  steam  formed 
by  the  combustion  of  the  hydrogen  is  condensed,  the  water  collecting 
in  drops  on  the  sides  of  the  jar. 


FIG. 12 


HYDROGEN 


35 


Precautions.  In  order  to  avoid  danger  it  is  absolutely  necessary  to 
prove  that  the  hydrogen  is  free  from  air  before  igniting  it.  This  can 
be  done  by  testing  small  amounts  of  the  escaping  gas.  A  convenient 
and  safe  method  of  doing  this  is  to  fill  a  test  tube  with  the  gas  by 
inverting  it  over  the  jet.  The  hydrogen,  on  account  of  its  lightness, 
collects  in  the  tube,  displacing  the  air.  After  holding  it  over  the  jet 
for  a  few  moments  in  order  that  it  may  be  filled  with  the  gas,  the  tube 
is  gently  brought,  mouth  downward,  to  the  flame  of  a  burner  placed 
not  nearer  than  an  arm's  length  from  the  jet.  If  the  hydrogen  is  mixed 
with  air  a  slight  explosion  occurs,  but  if  pure  it  burns  quietly  in  the 
tube.  The  operation  is  repeated  until  the  gas  burns  quietly,  when  the 
tube  is  quickly  brought  back  over  the  jet 
for  an  instant,  whereby  the  escaping  hydro- 
gen is  ignited  by  the  flame  in  the  tube. 

A  mixture  of  hydrogen  and  oxygen  is 
explosive.  That  a  mixture  of  hydrogen 
and  air  is  explosive  may  be  shown  safely 
as  follows :  A  cork  through  which  passes 
a  short  glass  tube  about  i  cm.  in  diameter 
is  fitted  air-tight  into  the  tubule  of  a  bell 
jar  of  2  1.  or  3  1.  capacity.  (A  thick  glass 
bottle  with  bottom  removed  may  be  used.) 
The  tube  is  closed  with  a  small  rubber 
stopper  and  the  bell  jar  filled  with  hydro- 
gen, the  gas  being  collected  over  water. 


FIG.  13 


When  entirely  filled  with  the  gas  the  jar  is  removed  from  the  water 
and  supported  by  blocks  of  wood  in  order  to  leave  the  bottom  of  the 
jar  open,  as  shown  in  Fig.  13.  The  stopper  is  now  removed  from  the 
tube  in  the  cork,  and  the  hydrogen,  which  on  account  of  its  lightness 
escapes  from  the  tube,  is  at  once  lighted.  As  the  hydrogen  escapes, 
the  air  flows  in  at  the  bottom  of  the  jar  and  mixes  with  the  remain- 
ing portion  of  the  hydrogen,  so  that  a  mixture  of  the  two  soon  forms, 
and  a  loud  explosion  results.  The  explosion  is  not  dangerous,  since 
the  bottom  of  the  jar  is  open,  thus  leaving  room  for  the  expansion  of 
the  hot  gas. 

Since  air  is  only  one  fifth  oxygen,  the  remainder  being  inert  gases, 
it  may  readily  be  inferred  that  a  mixture  of  hydrogen  with  pure  oxygen 
would  be  far  more  explosive  than  a  mixture  of  hydrogen  with  air. 
Such  mixtures  should  not  be  made  except  in  small  quantities  and  by 
experienced  workers. 


36       AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


Hydrogen  does  not  support  combustion.  While  hydrogen 
is  readily  combustible,  it  is  not  a  supporter  of  combustion. 
In  other  words,  substances  will  not  burn  in  it. 
This  may  be  shown  by  bringing  a  lighted  candle 
supported  by  a  stiff  wire  into  a  bottle  or  cylinder 
of  the  pure  gas,  as  shown  in  Fig.  14.  The  hydro- 
gen is  ignited  by  the  flame  of  the  candle  and 
burns  at  the  mouth  of  the  bottle,  where  it  comes 
in  contact  with  the  oxygen  in  the  air.  When  the 
candle  is  thrust  up  into  the  gas,  its  flame  is  ex- 
tinguished on  account  of  the  absence  of  oxygen. 
If  slowly  withdrawn,  the  candle  is  relighted  as  it 
passes  through  the  layer  of  burning  hydrogen. 

Reduction.    On  account  of  its  great  affinity  for 
oxygen,  hydrogen  has  the  power  of  abstracting  it  from 


O 


FIG. 14 


B 


FIG.  15 


many  of  its  compounds.    Thus,  if  a   stream  of  hydrogen, 
dried  by  passing  through  the  tube  B  (Fig.  15),  filled  with 


HYDROGEN  37 

calcium  chloride,  is  conducted  through  the  tube  C  contain- 
ing some  copper  oxide,  heated  to  a  moderate  temperature, 
the  hydrogen  abstracts  the  oxygen  from  the  copper  oxide. 
The  change  may  be  represented  as  follows : 

f  copper  1  £?       f  hydrogen  1  ^ 

hydrogen  +  J  f  II  =  1  f  *  +  copper 

L  oxygen  J  £  °        [  oxygen      J  *, 

The  water  formed  collects  in  the  cold  portions  of  the  tube  C 
near  its  end.  In  this  experiment  the  copper  oxide  is  said  to 
undergo  reduction.  Reduction  may  therefore  be  defined  as 
the  process  of  withdrawing  oxygen  from  a  compound. 

Relation  of  reduction  to  oxidation.  At  the  same  time 
that  the  copper  oxide  is  reduced  it  is  clear  that  the  hydro- 
gen is  oxidized,  for  it  combines  with  the  oxygen  given  up 
by  the  copper  oxide.  The  two  processes  are  therefore  very 
closely  related,  and  it  usually  happens  that  when  one  sub- 
stance is  oxidized  some  other  substance  is  reduced.  That 
substance  which  gives  up  its  oxygen  is  called  an  oxidizing 
agent,  while  the  substance  which  unites  with  the  oxygen  is 
called  a  reducing  agent. 

The  oxyhydrogen  blowpipe.  This  is  a  form  of  apparatus  used  for 
burning  hydrogen  in  pure  oxygen.  As  has  been  previously  stated,  the 
flame  produced 
by  the  combus- 
tion of  hydro-  C 
gen  in  the  air  is 
very  hot.  It  is 
evident  that  if 

pure  oxygen  is  ff  ^^ 

substituted  for  * 

air,  the  temperature  reached  will  be  much  higher,  since  there  are  no 
inert  gases  to  absorb  the  heat.  The  oxyhydrogen  blowpipe,  used  to 
effect  this  combination,  consists  of  a  small  tube  placed  within  a  larger 
one,  as  shown  in  Fig.  16. 


38       AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

The  hydrogen,  stored  under  pressure,  generally  in  steel  cylinders, 
is  first  passed  through  the  outer  tube  and  ignited  at  the  open  end  of 
the  tube.  The  oxygen  from  a  similar  cylinder  is  then  conducted 
through  the  inner  tube,  and  mixes  with  the  hydrogen  at  the  end  of  the 
tube.  In  order  to  produce  the  maximum  heat,  the  hydrogen  arid  oxy- 
gen must  be  admitted  to  the  blowpipe  in  the  exact  proportion  in  which 
they  combine,  viz.,  2  volumes  of  hydrogen  to  I  of  oxygen,  or  by  weight, 
i  part  of  hydrogen  to  7.94  parts  of  oxygen.  The  intensity  of  the  heat 
may  be  shown  by  bringing  into  the  flame  pieces  of  metal  such  as  iron 
wire  or  zinc.  These  burn  with  great  brilliancy.  Even  platinum,  having 
a  melting  point  of  1779°,  may  be  melted  by  the  heat  of  the  flame. 

While  the  oxyhydrogen  flame  is  intensely  hot,  it  is  almost  non- 
luminous.  If  directed  against  some  infusible  substance  like  ordinary 

lime  (calcium  oxide),  the  heat  is  so 
intense  that  the  lime  becomes  incan- 
descent and  glows  with  a  brilliant  light. 
This  is  sometimes  used  as  a  source  of 
light,  under  the  name  of  Drummond  or 

lime  light. 
-  GQ,S 

The  blast  lamp.    A  similar  form  of 

apparatus  is  commonly  used  in  the 
laboratory  as  a  source  of  heat  under 
the  name  blast  lamp  (Fig.  17).  This 
differs  from  the  oxyhydrogen  blowpipe 
only  in  the  size  of  the  tubes.  In  place 
of  the  hydrogen  and  oxygen  the  more 

accessible  coal  gas  and  air  are  respectively  used.  The  former  is  com- 
posed largely  of  a  mixture  of  free  hydrogen  and  gaseous  compounds 
of  carbon  and  hydrogen.  While  the  temperature  of  the  flame  is  not  so 
high  as  that  of  the  oxyhydrogen  blowpipe,  it  nevertheless  suffices  for 
most  chemical  operations  carried  out  in  the  laboratory. 

Uses  of  hydrogen.  On  account  of  its  cost,  hydrogen  is 
but  little  used  for  commercial  purposes.  It  is  sometimes 
used  as  a  material  for  the  inflation  of  balloons,  but  usually 
the  much  cheaper  coal  gas  is  substituted  for  it.  Even  hot 
air  is  often  used  when  the  duration  of  ascension  is  very 
short.  It  has  been  used  also  as  a  source  of  heat  and  light 


HYDROGEN 


39 


in  the  oxy hydrogen  blowpipe.  Where  the  electric  current 
is  available^  however,  this  form  of  apparatus  has  been  dis- 
placed almost  entirely  by  the  electric  light  and  electric  fur- 
nace, which  are  much  more  economical  and  more  powerful 
sources  of  light  and  heat. 

EXERCISES 

1.  Will  a  definite  weight  of  iron  decompose  an  unlimited  weight 
of  steam  ? 

2.  Why  is  oxygen  passed  through  the  inner  tube  of  the  oxy  hydro- 
gen blowpipe  rather  than  the  outer? 

3.  In  Fig.  14,  will  the  flame  remain  at  the  mouth  of  the  tube? 

4.  From  Fig.  15,  suggest  a  way  for  determining  experimentally 
the  quantity  of  water  formed  in  the  reaction. 

5.  Distinguish  clearly  between  the  following  terms  :  oxidation, 
reduction,  combustion,  and  kindling  temperature. 

6.  Is  oxidation  always  accompanied  by  reduction  ? 

7.  What  is  the  source  of  heat  in  the  lime  light  ?    What  is  the 
exact  use  of  lime  in  this  instrument  ? 

8.  In  Fig.  1 2,  why  is  it  necessary  to  dry  the  hydrogen  by  means 
of  the  calcium  chloride  in  the  tube  X? 

9.  At  what  pressure  would   the  weight  of  I  1.  of  hydrogen  be 
equal  to  that  of  oxygen  under  standard  conditions  ? 

10.  (a)  What  weight  of  hydrogen  can  be  obtained  from  1 50  g.  of 
sulphuric  acid  ?     (b}  What  volume  would  this  occupy  under  standard 
conditions?   (c}  The  density  of  sulphuric  acid  is  1.84.  What  volume 
would  the  1 50  g.  of  the  acid  occupy  ? 

11.  How  many  liters  of  hydrogen  can  be  obtained  from  50  cc.  of 
sulphuric  acid  having  a  density  of  1.84? 

12.  Suppose  you  wish  to  fill  five  liter  bottles  with  hydrogen,  the  gas 
to  be  collected  over  water  in  your  laboratory,  how  many  cubic  centi- 
meters of  sulphuric  acid  would  be  required  ? 


CHAPTER   IV 

COMPOUNDS  OF  HYDROGEN  AND  OXYGEN;    WATER  AND 
HYDROGEN  DIOXIDE 

WATER 

Historical.  Water  was  long  regarded  as  an  element.  In 
1781  Cavendish  showed  that  it  is  formed  by  the  union  of 
hydrogen  and  oxygen.  Being  a  believer  in  the  phlogiston 
theory,  however,  he  failed  to  interpret  his  results  correctly  A 
A  few  years  later  Lavoisier  repeated  Cavendish's  experi- 
ments and  showed  that  water  must  be  regarded  as  a  com- 
pound of  hydrogen  and  oxygen. 

General  methods  employed  for  the  determination  of  the  com- 
position of  a  compound.  The  composition  of  a  compound 
may  be  determined  by  either  of  two  general  processes ; 
these  are  known  as  analysis  and  synthesis. 

1.  Analysis  is  the  process  of  decomposing  a  compound 
into  its  constituents  and  determining  what  these  constituents 
are.    The  analysis  is  qualitative  when  it  results  in  merely 
determining  what  elements  compose  the  compound ;    it  is 
quantitative  when  the  exact  percentage  of  each  constituent 
is  determined.    Qualitative  analysis  must  therefore  precede 
quantitative  analysis,  for  it  must  be  known  what  elements 
are  in  a  compound  before  a  method  can  be  devised  for 
determining  exactly  how  much  of  each  is  present. 

2.  Synthesis  is  the  process  of  forming  a  compound  from 
its  constituent  parts.    It  is  therefore  the  reverse  of  analysis. 
Like  analysis,  it  may  be  either  qualitative  or  quantitative. 

40 


COMPOUNDS  OF  HYDROGEN  AND  OXYGEN 


Application  of  these  methods  to  the  determination  of  the 
composition  of  water.  The  determination  of  the  composition 
of  water  is  a  matter  of  great  interest  not  only  because  of  the 
importance  of  the  compound  but  also  because  the  methods 
employed  illustrate  the  general  methods  of  analysis  and 
synthesis. 

Methods  based  on  analysis.  The  methods  based  on  analy- 
sis may  be  either  qualitative  or  quantitative  in  character, 
i.  Qualitative  analysis.  As  was  stated  in  the  study  of 
oxygen,  water  may  be  separated  into  its  component  parts  by 
means  of  the  electric  current.  The  form 
of  apparatus  ordinarily  used  for  effecting 
this  analysis  is  shown  in  Fig.  18.  A  plati- 
num wire,  to  the  end  of  which  is  attached 
a  small  piece  of  platinum  foil  (about  1 5  mm. 
by  25  mm.),  is  fused  through  each  of  the 
tubes  B  and  D,  as  shown  in  the  figure. 
The  stopcocks  at  the  ends  of  these  tubes 
are  opened  and  water,  to  which  has  been 
added  about  one  tenth  of  its  volume  of 
sulphuric  acid,  is  poured  into 
the  tube  A  until  the  side  tubes 
B  and  D  are  completely  filled. 
The  stopcocks  are  then  closed. 
The  platinum  wires  extend- 
ing into  the  tubes  B  and  D  are  now  connected  with  the 
wires  leading  from  two  or  three  dichromate  cells  joined 
in  series.  The  pieces  of  platinum  foil  within  the  tubes  thus 
become  the  electrodes,  and  the  current  flows  from  one  to 
the  other  through  the  acidulated  water.  As  soon  as  the 
current  passes,  bubbles  of  gas  rise  from  each  of  the  elec- 
trodes and  collect  in  the  upper  part  of  the  tubes.  The  gas 


FIG. 18 


42       AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


rising  from  the  negative  electrode  is  found  to  be  hydrogen, 
while  that  from  the  positive  electrode  is  oxygen.  It  will  be 
seen  that  the  volume  of  the  hydrogen  is  approximately 
double  that  of  the  oxygen.  Oxygen  is  more  soluble  in 
water  than  hydrogen,  and  a  very  little  of  it  is  also  lost  by 
being  converted  into  ozone  and  other  substances.  It  has 

been  found  that  when  the 
necessary  corrections  are 
made  for  the  error  due  to 
these  facts,  the  volume  of  the 
hydrogen  is  exactly  double 
that  of  the  oxygen. 

Fig.  19  illustrates  a  simpler 
form  of  apparatus,  which  may 
be  used  in  place  of  that  shown 
in  Fig.  1 8.  A  glass  or  porce- 

lain  dish  is  Partially  mied 

FlG-  J9  with  water  to  which  has  been 

added  the  proper  amount  of  acid.  Two  tubes  filled  with  the 
same  liquid  are  inverted  over  the  electrodes.  The  gases 
resulting  from  the  decomposition  of  the  water  collect  in 
the  tubes. 

2.  Quantitative  analysis.  The  analysis  just  described 
is  purely  qualitative  and  simply  shows  that  water  contains 
hydrogen  and  oxygen.  It  does  not  prove  the  absence  of 
other  elements  ;  indeed  it  does  not  prove  that  the  hydrogen 
and  oxygen  are  present  in  the  proportion  in  which  they  are 
liberated  by  the  electric  current.  The  method  may  be  made 
quantitative,  however,  by  weighing  the  water  decomposed  and 
also  the  hydrogen  and  oxygen  obtained  in  its  decomposition. 
If  the  combined  weights  of  the  hydrogen  and  oxygen  exactly 
equal  the  weight  of  the  water  decomposed,  then  it  would 


COMPOUNDS  OF  HYDROGEN  AND  OXYGEN   43 

be  proved  that  the  water  consists  of  hydrogen  and  oxygen 
in  the  proportion  in  which  they  are  liberated  by  the  electric 
current.  This  experiment  is  difficult  to  carry  out,  however, 
so  that  the  more  accurate  methods  based  on  synthesis  are 
used. 

Methods  based  on  synthesis.    Two  steps  are  necessary 
to  ascertain  the  exact  composition  of  water  by  synthesis : 

(1)  to  show  by  qualitative  synthesis  that  water  is  formed 
by  the  union  of  oxygen  with  hydrogen  ; 

(2)  to  determine  by  quantitative  synthe- 
sis in  what  proportion  the  two  elements 
unite  to  form  water.    The  fact  that  water 
is  formed  by  the  combination  of  oxygen 
with  hydrogen  was  proved  in  the  preced- 
ing chapter.    The  quantitative  synthesis 
may  be  made  as  follows  : 

The  combination  of  the  two  gases  is 
brought  about  in  a  tube  called  a  eudi- 
ometer. This  is  a  graduated  tube  about 
60  cm.  long  and  2  cm.  wide,  closed  at 

one  end  (Fig.  20).    Near  the  closed  end 

r         .     .  FIG.  20 

two  platinum  wires  are  fused  through  the 

glass,  the  ends  of  the  wires  within  the  tube  being  separated 
by  a  space  of  2  mm.  or  3  mm.  The  tube  is  entirely  filled 
with  mercury  and  inverted  in  a  vessel  of  *  the  same  liquid. 
Pure  hydrogen  is  passed  into  the  tube  until  it  is  about  one 
fourth  filled.  The  volume  of  the  gas  is  then  read  off  on  the 
scale  and  reduced  to  standard  conditions.  Approximately 
an  equal  volume  of  pure  oxygen  is  then  introduced  and  the 
volume  again  read  off  and  reduced  to  standard  conditions. 
This  I  e  of  the  two  gases.  From  this 

the  ^  introduced  may  be  determined  by 


44       AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

subtracting  from  it  the  volume  of  the  hydrogen.  The  com- 
bination of  the  two  gases  is  now  brought  about  by  connect- 
ing the  two  platinum  wires  with  an  induction  coil  and 
passing  a  spark  from  one  wire  to  the  other.  Immediately  a 
slight  explosion  occurs.  The  mercury  in  the  tube  is  at  first 
depressed  because  of  the  expansion  of  the  gases  due  to  the 
heat  generated,  but  at  once  rebounds,  taking  the  place  of 
the  gases  which  have  combined  to  form  water.  The  volume 
of  the  water  in  the  liquid  state  is  so  small  that  it  may  be 
disregarded  in  the  calculations.  In  order  that  the  tempera- 
ture of  the  residual  gas  and  the  mercury  may  become  uni- 
form, the  apparatus  is  allowed  to  stand  for  a  few  minutes. 
The  volume  of  the  gas  is  then  read  off  and  reduced  to 
standard  conditions,  so  that  it  may  be  compared  with  the 
volumes  of  the  hydrogen  and  oxygen  originally  taken.  The 
residual  gas  is  then  tested  in  order  to  ascertain  whether  it 
is  hydrogen  or  oxygen,  experiments  having  proved  that  it  is 
never  a  mixture  of  the  two.  From  the  information  thus 
obtained  the  composition  of  the  water  may  be  calculated. 
Thus,  suppose  the  readings  were  as  follows  : 

Volume  of  hydrogen  taken 20.3  cc. 

Volume  of  hydrogen  and  oxygen 38.7 

Volume  of  oxygen 18.4 

Volume  of  gas  left  after  combination  has  taken 

place  (oxygen) 8.3 

The  20.3  cc.  of  hydrogen  ha  [  cc. 

minus  8.3  cc.  (or  10.1  cc.)  of  c  rcly  2 

volumes  of  hydrogen  have   combi  vygen. 

Since  oxygen  is  15.88   times  as  the 

proportion  by  weight  in  whic'i  •'.  -e  is 
I  part  of  hydrogen  to  7.94  of  o 


COMPOUNDS  OF  HYDROGEN  AND  OXYGEN  45 

Precaution.  If  the  two  gases  are  introduced  into  the  eudiometer  in 
the  exact  proportions  in  which  they  combine,  after  the  combination 
has  taken  place  the  liquid  will  rise  and  com- 
pletely fill  the  tube.  Under  these  conditions, 
however,  the  tube  is  very  likely  to  be  broken 
by  the  sudden  upward  rush  of  the  liquid. 
Hence  in  performing  the  experiment  care  is 
taken  to  introduce  an  excess  of  one  of  the 
gases. 

A  more  convenient  form  of  eudiometer.  A 
form  of  eudiometer  (Fig.  21 )  different  from  that 
shown  on  page  43  is  sometimes  used  to  avoid 
the  calculations  necessary  in  reducing  the  vol- 
umes of  the  gases  to  the  same  conditions  of 
temperature  and  pressure  in  order  to  make 
comparisons.  With  this  apparatus  it  is  pos- 
sible to  take  the  readings  of  the  volumes  under 
the  same  conditions  of  temperature  and  pres- 
sure, and  thus  compare  them  directly.  The 
apparatus  (Fig.  21)  is  filled  with  mercury  and 
the  gases  introduced  into  the  tube  A.  The 

experiment  is  carried  out  as  in  the  preceding  one,  except  that  before 
taking  the  reading  of  the  gas  volumes,  mercury  is  either  added  to  the 
tube  B  or  withdrawn  from  it  by  means  of  the  stopcock  C,  until  it  stands 
at  exactly  the  same  height  in  both  tubes.  The  gas  inclosed  in  tube  A 
is  then  under  atmospheric  pressure  ;  and  since  but  a  few  minutes  are 
required  for  performing  the  experiment,  the  conditions  of  temperature 
and  pressure  may  be  regarded  as  constant.  Hence  the  volumes  of  the 
hydrogen  and  oxygen  and  of  the  residual  gas  may  be  read  off  from 
the  tube  and  directly  compared. 

Method  used  by  Berzelius  and  Dumas.  The  method 
used  by  these  investigators  enables  us  to  determine  directly 
the  proportion  by  weight  in  which  the  hydrogen  and  oxy- 
gen combine.  Fig.  22  illustrates  the  apparatus  used  in 
making  this  determination.  B  is  a  glass  tube  containing 
copper  oxide.  C  and  D  are  glass  tubes  filled  with  calcium 
chloride,  a  substance  which  has  great  affinity  for  water. 


46       AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


The  tubes  B  and  C,  including  their'  contents,  are  carefully 
weighed,  and  the  apparatus  connected  as  shown  in  the 
figure.  A  slow  current  of  pure  hydrogen  is  then  passed 
through  A,  and  that  part  of  the  tube  B  which  contains 
copper  oxide  is  carefully  heated.  The  hydrogen  combines 
with  the  oxygen  present  in  the  copper  oxide  to  form  water, 
which  is  absorbed  by  the  calcium  chloride  in  tube  C.  The 
calcium  chloride  in  tube  D  prevents  any  moisture  entering 
tube  C  from  the  air.  The  operation  is  continued  until  an 

B  a 

~n=r 


FIG.  22 

appreciable  amount  of  water  has  been  formed.  The  tubes 
B  and  C  are  then  weighed  once  more.  The  loss  of  weight 
in  the  tube  B  will  exactly  equal  the  weight  of  oxygen  taken 
up  from  the  copper  oxide  in  the  formation  of  the  water. 
The  gain  in  weight  in  the  tube  C  will  exactly  equal  the 
weight  of  the  water  formed.  The  difference  in  these  weights 
will  of  course  equal  the  weight  of  the  hydrogen  present  in 
the  water  formed. 

Dumas'  results.  The  above  method  for  the  determina- 
tion of  the  composition  of  water  was  first  used  by  Berzelius 
in  1820.  The  work  was  repeated  in  1843  by  Dumas,  the 
average  of  whose  results  is  as  follows  : 

Weight  of  water  formed 236.36  g. 

Oxygen  given  up  by  the  copper  oxide     .     .     210.04 

Weight  of  hydrogen  present  in  water     .       26.32 


COMPOUNDS  OF  HYDROGEN  AND  OXYGEN      47 

According  to  this  experiment  the  ratio  of  hydrogen  to  oxy- 
gen in  water  is  therefore  26.32  to  210.04,  or  as  I  to  7-98- 

Morley 's  results.  The  American  chemist  Morley  has 
recently  determined  the  composition  of  water,  extreme  pre- 
cautions being  taken  to  use  pure  materials  and  to  elimi- 
nate all  sources  of  error.  The  hydrogen  and  oxygen  which 
combined,  as  well  as  the  water  formed,  were  all  accurately 
weighed.  According  to  Morley 's  results,  i  part  of  hydrogen 
by  weight  combines  with  7.94  parts  of  oxygen  to  form  water. 

Comparison  of  results  obtained.  From  the  above  discus- 
sions it  is  easy  to  see  that  it  is  by  experiment  alone  that 
the  composition  of  a  compound  can  be  determined.  Differ- 
ent methods  may  lead  to  slightly  different  results.  The 
more  accurate  the  method  chosen  and  the  greater  the  skill 
with  which  the  experiment  is  carried  out,  the  more  accurate 
will  be  the  results.  It  is  generally  conceded  by  chemists 
that  the  results  obtained  by  Morley  in  reference  to  the 
composition  of  water  are  the  most  accurate  ones.  In  accord- 
ance with  these  results,  then,  water  must  be  regarded  as 
a  compound  containing  hydrogen  and  oxygen  in  the  propor- 
tion of  /  part  by  weight  of  hydrogen  to  7.94 parts  by  weight 
of  oxygen. 

Relation  between  the  volume  of  aqueous  vapor  and  the 
volumes  of  the  hydrogen  and  oxygen  which  combine  to 
form  it.  When  the  quantitative  synthesis  of  water  is  car- 
ried out  in  the  eudiometer  as  described  above,  the  water 
vapor  formed  by  the  union  of  the  hydrogen  and  oxygen  at 
once  condenses.  The  volume  of  the  resulting  liquid  is  so 
small  that  it  may  be  disregarded  in  making  the  calculations. 
If,  however,  the  experiment  is  carried  out  at  a  temperature 
of  1 00°  or  above,  the  water  vapor  formed  is  not  condensed 
and  it  thus  becomes  possible  to  compare  the  volume  of  the 


48        AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

vapor  with  the  volumes  of  hydrogen  and  oxygen  which 
combined  to  form  it.  This  can  be  accomplished  by  sur- 
rounding the  arm  A  of  the  eudiometer  (Fig.  23)  with  the 
tube  B  through  which  is  passed  the  vapor  obtained  by 
boiling  some  liquid  which  has  a  boiling  point  above  100°. 
In  this  way  it  has  been  proved  that  2  volumes  of  hydro- 
gen and  i  volume  of  oxygen  combine  to  form  exactly  2  vol- 
umes of  water  vapor,  the  volumes  all  being  measured  under 
the  same  conditions  of  temperature  and 
pressure.  It  will  be  noted  that  the  re- 
lation between  these  volumes  may  be 
expressed  by  whole  numbers.  The  sig- 
nificance of  this  very  important  fact 
will  be  discussed  in  a  subsequent 
chapter. 

Occurrence  of  water.    Water  not  only 
covers  about  three  fourths  of  the  surface 
of  the  earth,  and  is  present  in  the  atmos- 
phere in  the  form  of  moisture,  but  it  is 
FIG  2^  also  a  common  constituent  of  the  soil 

and  rocks  and  of  almost  every  form  of 
animal  and  vegetable  organism .  The  human  body  is  nearly 
70%  water.  This  is  derived  not  only  from  the  water  which 
we  drink  but  also  from  the  food  which  we  eat,  most  of 
which  contains  a  large  percentage  of  water.  Thus  potatoes 
contain  about  78%  of  water,  milk  85%,  beef  over  50%, 
apples  84%,  tomatoes  94%. 

Impurities  in  water.  Chemically  pure  water  contains 
only  hydrogen  and  oxygen.  Such  a  water  never  occurs  in 
nature,  however,  for  being  a  good  solvent,  it  takes  up 
certain  substances  from  the  rocks  and  soil  with  which  it 
comes  in  contact.  When  such  waters  are  evaporated  these 


COMPOUNDS  OF  HYDROGEN  AND  OXYGEN      49 

substances  are  deposited  in  the  form  of  a  residue.  Even  rain 
water,  which  is  the  purest  form  occurring  in  nature,  contains 
dust  particles  and  gases  dissolved  from  the  atmosphere. 
The  foreign  matter  in  water  is  of  two  kinds,  namely,  mineral, 
such  as  common  salt  and  limestone,  and  organic,  that  is  the 
products  of  animal  -and  vegetable  life. 

Mineral  matter  in  water.  The  amount  and  nature  of  the  mineral 
matter  present  in  different  waters  vary  greatly,  depending  on  the 
character  of  the  rocks  and  soil  with  which  the  waters  come  in  contact. 
The  more  common  of  the  substances  present  are  common  salt  and 
compounds  of  calcium,  magnesium,  and  iron.  One  liter  of  the  average 
river  water  contains  about  175  mg.  of  mineral  matter.  Water  from 
deep  wells  naturally  contains  more  mineral  matter  than  river  water, 
generally  two  or  three  times  as  much,  while  sea  water  contains  as 
much  as  35,000  mg.  to  the  liter. 

Effect  of  impurities  on  health.  The  mineral  matter  in 
water  does  not,  save  in  very  exceptional  cases,  render  the 
water  injurious  to  the  human  system.  In  fact  the  presence 
of  a  certain  amount  of  such  matter  is  advantageous,  supply- 
ing the  mineral  constituents  necessary  for  the  formation 
of  the  solid  tissues  of  the  body.  The  presence  of  organic 
matter,  on  the  other  hand,  must  always  be  regarded  with 
suspicion.  This  organic  matter  may  consist  not  only  of  the 
products  of  animal  and  vegetable  life  but  also  of  certain 
microscopic  forms  of  living  organisms  which  are  likely  to 
accompany  such  products.  Contagious  diseases  are  known 
to  be  due  to  the  presence  in  the  body  of  minute  living 
organisms  or  germs.  Each  disease  is  caused  by  its  own 
particular  kind  of  germ.  Through  sewage  these  germs  may 
find  their  way  from  persons  afflicted  with  disease  into  the 
water  supply,  and  it  is  principally  through  the  drinking  water 
that  certain  of  these  diseases,  especially  typhoid  fever,  are 
spread.  It  becomes  of  great  importance,  therefore,  to  be 


50       AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

able  to  detect  such  matter  when  present  in  drinking  water 
as  well  as  to  devise  methods  whereby  it  can  be  removed  or 
at  least  rendered  harmless. 

Analysis  of  water.  The  mineral  analysis  of  a  water  is,  as  the  name 
suggests,  simply  the  determination  of  the  mineral  matter  present. 
Sanitary  analysis,  on  the  other  hand,  is  the  determination  of  the 
organic  matter  present.  The  physical  properties  of  a  water  give  no 
conclusive  evidence  as  to  its  purity,  since  a  water  may  be  unfit  for 
drinking  purposes  and  yet  be  perfectly  clear  and  odorless.  Neither 


FIG.  24 

can  any  reliance  be  placed  on  the  simple  methods  often  given  for  test- 
ing the  purity  of  water.  Only  the  trained  chemist  can  carry  out  such 
methods  of  analysis  as  can  be  relied  upon. 

Purification  of  water.  Three  general  methods  are  used 
for  the  purification  of  water,  namely,  distillation,  filtration, 
and  boiling. 

i.  Distillation.  The  most  effective  way  of  purifying 
natural  waters  is  by  the  process  of  distillation.  This  con- 
sists in  boiling  the  water  and  condensing  the  steam.  Fig.  24 
illustrates  the  process  of  distillation,  as  commonly  conducted 


COMPOUNDS  OF  HYDROGEN  AND  OXYGEN   51 

in  the  laboratory.  Ordinary  water  is  poured  into  the  flask 
A. and  boiled.  The  steam  is  conducted  through  the  con- 
denser B,  which  consists  essentially  of  a  narrow  glass  tube 
sealed  within  a  larger  one,  the  space  between  the  two  being 
filled  with  cold  water,  which  is  admitted  at  £Tand  escapes 
at  D.  The  inner  tube  is  thus  kept  cool  and  the  steam  in 
passing  through  it  is  condensed.  The  water  formed  by  the 
condensation  of  the  steam  collects  in  the  receiver  E  and  is 
known  as  distilled  water.  Such  water  is  practically  pure, 
since  the  impurities  are  nonvolatile  and  remain  in  the 
flask  A. 

Commercial  distillation.  In  preparing  distilled  water  on  a  large 
scale,  the  steam  is  generated  in  a  boiler  or  other  metal  container  and 
condensed  by  passing  it  through  a  pipe  made  of  metal,  generally  tin. 
This  pipe  is  wound  into  a  spiral  and  is  surrounded  by  a  current  of 
cold  water.  Distilled  water  is  used  by  the  chemist  in  almost  all  of  his 
work.  It  is  also  used  in  the  manufacture  of  artificial  ice  and  for 
drinking  water. 

Fractional  distillation.  In  preparing  distilled  water,  it  is  evident 
that  if  the  natural  water  contains  some  substance  which  is  volatile 
its  vapor  will  pass  over  and  be  condensed  with  the  steam,  so  that  the 
distillate  will  not  be  pure  water.  Even  such  mixtures,  however,  may 
generally  be  separated  by  repeated  distillation.  Thus,  if  a  mixture  of 
water  (boiling  point  100°)  and  alcohol  (boiling  point  78°)  is  distilled, 
the  alcohol,  having  the  lower  boiling  point,  tends  to  distill  first,  fol- 
lowed by  the  water.  The  separation  of  the  two  is  not  perfect,  how- 
ever, but  may  be  made  nearly  so  by  repeated  distillations.  The  process 
of  separating  a  mixture  of  volatile  substances  by  distillation  is  known 
as  fractional  distillation. 

2.  Filtration.  The  process  of  distillation  practically  re- 
moves all  nonvolatile  foreign  matter,  mineral  as  well  as 
organic.  In  purifying  water  for  drinking  purposes,  however, 
it  is  only  necessary  to  eliminate  the  latter  or  to  render  it 
harmless.  This  is  ordinarily  done  either  by  filtration  or 


52        AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


boiling.  In  filtration  the  water  is  passed  through  some 
medium  which  will  retain  the  organic  matter.  Ordinary 
charcoal  is  a  porous  substance  and  will  condense  within  its 
pores  the  organic  matter  in  water  if  brought  in  contact 
with  it.  It  is  therefore  well  adapted  to  the  construction  of 
filters.  Such  filters  to  be  effective  must  be  kept  clean,  since 
it  is  evident  that  the  charcoal  is  useless  after  its  pores  are 
filled.  A  more  effective  type  of  filter  is  the  Chamberlain- 
Pasteur  filter.  In  this  the  water  is  forced  through  a  porous 

cylindrical    cup, 


the  pores  being 
so  minute  as  to 
strain  out  the 
organic  matter. 

City  filtration 
beds.  For  purify- 
ing the  water  sup- 
ply of  cities,  large 
filtration  beds  are 
prepared  from  sand 
and  gravel,  and  the 
water  is  allowed  to 
filter  through  these. 


FIG. 25 


Some  of  the  impurities  are  strained  out  by  the  filter,  while  others  are 
decomposed  by  the  action  of  certain  kinds  of  bacteria  present  in  the 
sand.  Fig.  25  shows  a  cross  section  of  a  portion  of  the  filter  used  in 
purifying  the  water  supply  of  Philadelphia.  The  water  filters  through 
the  sand  and  gravel  and  passes  into  the  porous  pipe  A,  from  which 
it  is  pumped  into  the  city  mains.  The  filters  are  covered  to  prevent 
the  water  from  freezing  in  cold  weather. 

3.  Boiling.  A  simpler  and  equally  efficient  method  for 
purifying  water  for  drinking  purposes  consists  in  boiling  the 
water.  It  is  the  germs  in  water  that  render  it  dangerous  to 
health.  These  germs  are  living  forms  of  matter.  If  the 


COMPOUNDS  OF  HYDROGEN  AND  OXYGEN   53 

water  is  boiled,  the  germs  are  killed  and  the  water  rendered 
safe.  While  these  germs  are  destroyed  by  heat,  cold  has 
little  effect  upon  them.  Thus  Dewar,  in  working  with 
liquid  hydrogen,  exposed  some  of  these  minute  forms  of 
life  to  the  temperature  of  boiling  hydrogen  (—  252°)  with- 
out killing  them. 

Self-purification  of  water.  It  has  long  been  known  that 
water  contaminated  with  organic  matter  tends  to  purify 
itself  when  exposed  to  the  air.  This  is  due  to  the  fact  that 
the  water  takes  up  a  small  amount  of  oxygen  from  the  air, 
which  gradually  oxidizes  the  organic  matter  present  in  the 
water.  While  water  is  undoubtedly  purified  in  this  way, 
the  method  cannot  be  relied  upon  to  purify  a  contaminated 
water  so  as  to  render  it  "safe  for  drinking  purposes. 

Physical  properties.  Pure  water  is  an  odorless  and  taste- 
less liquid,  colorless  in  thin  layers,  but  having  a  bluish  tinge 
when  observed  through  a  considerable  thickness.  It  solidi- 
fies at  o°  and  boils  at  100°  under  the  normal  pressure  of 
one  atmosphere.  If  the  pressure  is  increased,  the  boiling 
point  is  raised.  When  water  is  cooled  it  steadily  contracts 
until  the  temperature  of  4°  is  reached :  it  then  expands. 
Water  is  remarkable  for  its  ability  to  dissolve  other  sub- 
stances, and  is  the  best  solvent  known.  Solutions  of  solids 
in  water  are  more  frequently  employed  in  chemical  work 
than  are  the  solid  substances,  for  chemical  action  between 
substances  goes  on  more  readily  when  they  are  in  solution 
than  it  does  when  they  are  in  the  solid  state. 

Chemical  properties.  Water  is  a  very  stable  substance, 
or,  in  other  words,  it  does  not  undergo  decomposition  readily. 
To  decompose  it  into  its  elements  by  heat  ^orle  requires  a 
very  high  temperature.  The  decomposition  begins  at  about 
1000°,  but  is  only  half  complete  at  2500°.  Though  very 


54        AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

stable  towards  heat,  water  can  be  decomposed  in  other  ways, 
as  by  t^e  action  of  the  electrical  current  or  by  certain 
metals. 

Heat  of  formation  and  heat  of  decomposition  are  equal.  The  fact 
that  a  very  high  temperature  is  necessary  to  decompose  water  into 
hydrogen  and  oxygen  is  in  accord  with  the  fact  that  a  great  deal  of 
heat  is  evolved  by  the  union  of  hydrogen  and  oxygen  ;  for  it  has  been 
proved  that  the  heat  necessary  to  decompose  a  compound  into  its 
elements  (heat  of  decomposition)  is  equal  to  the  heat  evolved  in  the 
formation  of  a  compound  from  its  elements  (heat  of  formation). 

Water  of  crystallization.  When  a  solid  is  dissolved  in 
water  and  the  resulting  solution  is  allowed  to  evaporate,  the 
solid  separates  out,  often  in  the  form  of  crystals.  It  has 
been  found  that  the  crystals  of  many  compounds,  although 
perfectly  dry,  give  up  a  definite  amount  of  water  when 
heated,  the  substance  at  the  same  time  losing  its  crystalline 
form.  Such  water  is  called  water  of  crystallization.  This 
varies  in  amount  with  different  compounds,  but  is  perfectly 
definite  for  the  same  compound.  Thus,  if  a  perfectly  dry 
crystal  of  copper  sulphate  is  strongly  heated  in  a  tube,  water 
is  evolved  and  jcondenses  on  the  sides  of  the  tube,  the  crys- 
tal crumbling  to  a  light  powder.  The  weight  of  the  water 
evolved  is  always  equal  to  exactly  36.07%  of  the  weight  of 
copper  sulphate  crystals  heated.  The  water  must  therefore 
be  in  chemical  combination  with  the  substance  composing 
the  crystal ;  for  if  simply  mixed  with  it  or  adhering  to  it, 
not  only  would  the  substance  appear  moist  but  the  amount 
present  would  undoubtedly  vary.  The  combination,  however, 
must  be  a  very  weak  one,  since  the  water  is  often  expelled 
by  even  a  gentle  heat.  Indeed,  in  some  cases  the  water  is 
given  up  on  simple  exposure  to  air.  Such  compounds  are 
said  to  be  efflorescent.  Thus  a  crystal  of  sodium  sulphate 


COMPOUNDS  OF  HYDROGEN  AND  OXYGEN   55 

(Glauber's  salt)  on  exposure  to  air  crumbles  to  a  fine  powder, 
owing  to  the  escape  of  its  water  of  crystallization.  Other 
substances  have  just  the  opposite  property :  they  absorb 
moisture  when  exposed  to  the  air.  For  example,  if  a  bit  of 
dry  calcium  chloride  is  placed  in  moist  air,  in  the  course  of 
a  few  hours  it  will  have  absorbed  sufficient  moisture  to  dis- 
solve it.  Such  substances  are  said  to  be  deliquescent.  A 
deliquescent  body  serves  as  a  good  drying  or  desiccating 
agent.  We  have  already  employed  calcium  chloride  as  an 
agent  for  absorbing  the  moisture  from  hydrogen.  Many 
substances,  as  for  example  quartz,  form  crystals  which  con- 
tain no  water  of  crystallization. 

Mechanically  inclosed  water.  Water  of  crystallization  must  be  care- 
fully distinguished  from  water  which,  is  mechanically  inclosed  in  a 
crystal  and  which  can  be  removed  by  powdering  the  crystal  and  dry- 
ing. Thus,  when  crystals  of  common  salt  are  heated,  the  water  inclosed 
in  the  crystal  is  changed  into  steam  and  bursts  the  crystal  with  a 
crackling  sound.  Such  crystals  are  said  to  decrepitate.  That  this 
water  is  not  combined  is  proved  by  the  fact  that  the  amount  present 
varies  and  that  it  has  all  the  properties  of  water. 

-Uses  of  water.  The  importance  of  water  in  its  relation 
to  life  and  commerce  is  too  well  known  to  require  comment. 
Its  importance  to  the  chemist  has  also  been  pointed  out. 
It  remains  to  call  attention  to  the  fact  that  it  is  used  as  a 
standard  in  many  physical  measurements-.  Thus  o°  and 
1 00°  on  the  centigrade  scale  are  respectively  the  freezing 
and  the  boiling  points  of  water  under  normal  pressure.  The 
weight  of  i  cc.  of  water  at  its  point  of  greatest  density  is 
the  unit  of  weight  in  the  metric  system,  namely,  the  gram. 
It  is  also  taken  as  the  unit  for  the  determination  of  the 
density  of  liquids  and  solids  as  well  as  for  the  measurement 
of  amounts  of  heat. 


56        AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


HYDROGEN  DIOXIDE 

Composition.  As  has  been  shown,  i  part  by  weight  of 
hydrogen  combines  with  7.94  parts  by  weight  of  oxygen  to 
form  water.  It  is  possible,  however,  to  obtain  a  second  com- 
pound of  hydrogen  and  oxygen  differing  from  water  in 
composition  in  that  I  part  by  weight  of  hydrogen  is  com- 
bined with  2  X  7.94,  or  15.88  parts,  of  oxygen.  This  com- 
pound is  called  hydrogen  dioxide  or  hydrogen  peroxide,  the 
prefixes  di-  and  per-  signifying  that  it  contains  more  oxy- 
gen than  hydrogen  oxide,  which  is  the  chemical  name  for 
water. 

Preparation.  Hydrogen  dioxide  cannot  be  prepared 
cheaply  by  the  direct  union  of  hydrogen  and  oxygen,  and 
indirect  methods  must  therefore  be  used.  It  is  commonly 
prepared  by  the  action  of  a  solution  of  sulphuric  acid  on 
barium  dioxide.  The  change  which  takes  place  may  be 
indicated  as  follows : 


* 

rs 

* 

0> 

r^ 

' 

hydrogen 
sulphur 
oxygen 

°o 

a 

.-C4-  « 
^  3  ~ 

1- 

3 

barium 
oxygen 

1 

'S  

"  e  ~~  " 

a 

H 

rt 

barium 
sulphur 
oxygen 

en 

Jo 

. 

-H 


hydrogen 
oxygen 


In  other  words,  the  barium  and  hydrogen  in  the  two  com- 
pounds exchange  places.  By  this  method  a  dilute  solution 
of  the  dioxide  in  water  is  obtained.  It  is  possible  to  sepa- 
rate the  dioxide  from  the  water  by  fractional  distillation. 
This  is  attended  with  great  difficulties,  however,  since  the 
pure  dioxide  is  explosive.  The  distillation  is  carried  on 
under  diminished  pressure  so  as  to  lower  the  boiling  points 
as  much  as  possible  ;  otherwise  the  high  temperature  would 
decompose  the  dioxide.  , 


COMPOUNDS  OF  HYDROGEN  AND  OXYGEN   57 

Properties-.  Pure  hydrogen  dioxide  is  a  colorless  sirupy 
liquid  having  a  density  of  1.49.  Its  most  characteristic 
property  is  the  ease  with  which  it  decomposes  into  water 
and  oxygen.  One  part  by  weight  of  hydrogen  is  capable  of 
holding  firmly  only  7.94  parts  of  oxygen.  The  additional  7.94 
parts  of  oxygen  present  in  hydrogen  dioxide  are  therefore 
easily  evolved,  the  compound  breaking  down  into  water  and 
oxygen.  This  decomposition  is  attended  by  the  generation 
of  considerable  heat.  In  dilute  solution  hydrogen  dioxide  is 
fairly  stable,  although  such  a  solution  should  be  kept  in  a 
dark,  cool  place,  since  both  heat  and  light  aid  in  the  decom- 
position of  the  dioxide. 

Uses.  Solutions  of  hydrogen  dioxide  are  used  largely  as 
oxidizing  agents.  The  solution  sold  by  druggists  contains 
3%  of  the  dioxide  and  is  used  in  medicine  as  an  anti- 
septic. Its  use  as  an  antiseptic  depends  upon  its  oxidizing 
properties. 

EXERCISES 

1.  Why  does  the  chemist  use  distilled  water  in  making  solutions, 
rather  than  filtered  water? 

2.  How  could  you  determine  the  total  amount  of  solid  matter 
dissolved  in  a  sample  of  water  ? 

3.  How  could  you  determine  whether  a  given  sample  of  water  is 
distilled  water  ? 

4.  How  could  the  presence  of  air  dissolved  in  water  be  detected  ? 

5.  How  could  the  amount  of  water  in  a  food  such  as  bread  or 
potato  be  determined  ? 

6.  Would  ice  frozen  from  impure  water  necessarily  be  free  from 
disease  germs  ? 

7.  Suppose  that  the  maximum  density  of  water  were  at  o°  in  place 
of  4°;  what  effect  would  this  have  on  the  formation  of  ice  on  bodies 
of  water  ? 

8.  Is  it  possible  for  a  substance  to  contain  both  mechanically 
inclosed  water  and  water  of  crystallization? 


58       AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

9.    If  steam  is  heated  to  2000°  and  again  cooled,  has  any  chem- 
ical change  taken  place  in  the  steam  ? 

10.  Why  is  cold  water  passed  into  C  instead  of  D  (Fig.  24)  ? 

11.  Mention  at  least  two  advantages  that  a  metal  condenser  has 
over  a  glass  condenser. 

12.  Draw  a  diagram  of  the  apparatus  used  in  your  laboratory  for 
supplying  distilled  water. 

13.  20  cc.  of  hydrogen  and  7  cc.  of  oxygen  are  placed  in  a  eudi- 
ometer and  the  mixture  exploded,     (a)  How  many  cubic  centimeters 
of  aqueous  vapor  are  formed  ?  (£)  What  gas   and  how  much  of  it 
remains  in  excess  ? 

14.  (a)  What  weight  of  water  can  be  formed  by  the  combustion 
of  100  1.  of  hydrogen,  measured  under  standard  conditions?  (<£)  What 
volume  of  oxygen  would  be  required  in(#)?     (V)  What  weight  of 
potassium  chlorate  is  necessary  to  prepare  this  amount  of  oxygen? 

15.  What  weight  of  oxygen  is  present  in  i  kg.  of  the  ordinary 
hydrogen  dioxide  solution?     In  the  decomposition  of  this  weight  of 
the  dioxide  into  water  and  oxygen,  what  volume  of  oxygen  (measured 
under  standard  conditions)  is  evolved  ? 


CHAPTER  V 
THE  ATOMIC  THEORY 

Three  fundamental  laws  of  matter.  Before  we  can  gain 
any  very  definite  idea  in  regard  to  the  structure  of  matter, 
and  the  way  in  which  different  kinds  of  substances  act 
chemically  upon  each  other,  it  is  necessary  to  have  clearly 
in  view  three  fundamental  laws  of  matter.  These  laws  have 
been  established  by  experiment,  and  any  conception  which 
may  be  formed  concerning  matter  must  therefore  be  in 
harmony  with  them.  The  laws  are  as  follows  : 

Law  of  conservation  of  matter.  This  law  has  already 
been  touched  upon  in  the  introductory  chapter,  and  needs 
no  further  discussion.  It  will  be  recalled  that  it  may  be 
stated  thus  :  Matter  can  neither  be  created  nor  destroyed, 
though  it  can  be  changed  from  one  form  into  another. 

Law  of  definite  composition.  In  the  earlier  days  of 
chemistry  there  was  much  discussion  as  to  whether  the 
composition  of  a  given  compound  is  always  precisely  the 
same  or  whether  it  is  subject  to  some  variation.  Two 
Frenchmen,  Berthollet  and  Proust,  were  the  leaders  in  this 
discussion,  and  a  great  deal  of  most  useful  experimenting 
was  done  to  decide  the  question.  Their  experiments,  as 
well  as  all  succeeding  ones,  have  shown  that  the  composi- 
tion of  a  pure  chemical  compound  is  always  exactly  the 
-same.  Water  obtained  by  melting  pure  ice,  condensing 
steam,  burning  hydrogen  in  oxygen,  has  always  11.18% 
hydrogen  and  88.82%  oxygen  in  it.  Red  oxide  of  mer- 
cury, from  whatever  source  it  is  obtained,  contains  92.6% 

59 


60       AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

mercury  and  7.4%  oxygen.  This  truth  is  known  as  the  law 
of  definite  composition,  and  may  be  stated  thus  :  The  com- 
position of  a  chemical  compound  never  varies. 

Law  of  multiple  proportion.  It  has  already  been  noted, 
however,  that  hydrogen  and  oxygen  combine  in  two  differ- 
ent ratios  to  form  water  and  hydrogen  dioxide  respectively. 
It  will  be  observed  that  this  fact  does  not  contradict  the 
law  of  definite  composition,  for  entirely  different  substances 
are  formed.  These  compounds  differ  from  each  other  in 
composition,  but  the  composition  of  each  one  is  always 
constant.  This  ability  of  two  elements  to  unite  in  more 
than  one  ratio  is  very  frequently  observed.  Carbon  and 
oxygen  combine  in  two  different  ratios ;  nitrogen  and  oxy- 
gen combine  to  form  as  many  as  five  distinct  compounds, 
each  with  its  own  precise  composition. 

In  the  first  decade  of  the  last  century  John  Dalton,  an 
English  school-teacher  and  philosopher,  endeavored  to  find 
some  rule  which  holds  between  the  ratios  in  which  two 
given  substances  combine.  His  studies  brought  to  light  a 
very  simple  relation,  which  the  following  examples  will  make 
clear.  In  water  the  hydrogen  and  oxygen  are  combined  in 
the  ratio  of  I  part  by  weight  of  hydrogen  to  7.94  parts  by 
weight  of  oxygen.  In  hydrogen  dioxide  the  I  part  by 
weight  of  hydrogen  is  combined  with  15.88  parts  by  weight 
of  oxygen.  The  ratio  between  the  amounts  of  oxygen 
which  combine  with  the  same  amount  of  hydrogen  to 
form  water  and  hydrogen  dioxide  respectively  is  therefore 
7.94:  15.88,  or  i  :  2. 

Similarly,  the  element  iron  combines  with  oxygen  to 
form  two  oxides,  one  of  which  is  black  and  the  other  red. 
By  analysis  it  has  been  shown  that  the  former  contains  I 
part  by  weight  of  iron  combined  with  0.286  parts  by  weight 


THE  ATOMIC  THEORY  6 1 

of  oxygen,  while  the  latter  contains  i  part  by  weight  of 
iron  combined  with  0.429  parts  by  weight  of  oxygen.  Here 
again  we  find  that  the  amounts  of  oxygen  which  combine 
with  the  same  fixed  amount  of  iron  to  form  the  two  com- 
pounds are  in  the  ratio  of  small  whole  numbers,  viz.,  2:3. 

Many  other  examples  of  this  simple  relation  might  be 
given,  since  it  has  been  found  to  hold  true  in  all  cases 
where  more  than  one  compound  is  formed  from  the  same 
elements.  Dalton's  law  of  multiple  proportion  states  these 
facts  as  follows  :  When  any  two  elements,  A  and  B,  combine 
to  form  more  than  one  compound,  the  amounts  of  B  which 
unite  with  any  fixed  amount  of  A  bear  the  ratio  of  small 
whole  numbers  to  each  other. 

Hypothesis  necessary  to  explain  the  laws  of  matter. 
These  three  generalizations  are  called  laws,  because  they 
express  in  concise  language  truths  which  are  found  by  care- 
ful experiment  to  hold  good  in  all  cases.  They  do  not  offer 
any  explanation  of  the  facts,  but  merely  state  them.  The 
human  mind,  however,  does  not  rest  content  with  the  mere 
bare  facts,  but  seeks  ever  to  learn  the  explanation  of  the 
facts.  A  suggestion  which  is  offered  to  explain  such  a  set 
of  facts  is  called  an  hypothesis.  The  suggestion  which 
Dalton  offered  to  explain  the  three  laws  of  matter,  called 
the  atomic  hypothesis,  was  prompted  by  his  view  of  the  con- 
stitution of  matter,  and  it  involves  three  distinct  assump- 
tions in  regard  to  the  nature  of  matter  and  chemical  action. 
Dalton  could  not  prove  these  assumptions  to  be  true,  but 
he  saw  that  if  they  were  true  the  laws  of  matter  become 
very  easy  to  understand. 

Dalton's  atomic  hypothesis.  The  three  assumptions  which 
Dalton  made  in  regard  to  the  nature  of  matter,  and  which 
together  constitute  the  atomic  hypothesis,  are  these : 


62        AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

1.  All   elements  are  made  up  of   minute,  independent 
particles  which  Dalton  designated  as  atoms. 

2.  All  atoms  of  the  same  element  have  equal  masses; 
those  of  different  elements  have  different  masses  ;  in  any 
change  to  which  an  atom  is  subjected  its  mass  does  not 
change. 

3.  When  two  or  more  elements  unite  to  form  a  com- 
pound,  the    action    consists    in    the    union  of    a    definite 
small  number  of  atoms  of  each  element  to  form  a  small 
particle  of    the    compound.    The  smallest    particles  of  a 
given  compound  are  therefore  exactly  alike  in  the  number 
and  kinds  of  atoms  which  they  contain,  and  larger  masses 
of  the  substances  are  simply  aggregations  of  these  least 
particles. 

Molecules  and  atoms.  Dalton  applied  the  name  atom  not 
only  to  the  minute  particles  of  the  elements  but  also  to  the 
least  particles  of  compounds.  Later  Avogadro,  an  Italian 
scientist,  pointed  out  the  fact  that  the  two  are  different, 
since  the  smallest  particle  of  an  element  is  a  unit,  while  that 
of  a  compound  must  have  at  least  two  units  in  it.  He  sug- 
gested the  name  molecule  for  the  least  particle  of  a  com- 
pound which  can  exist,  retaining  the  name  atom  for  the 
smallest  particle  of  an  element.  In  accordance  with  this 
distinction,  we  may  define  the  atom  and  the  molecule  as 
follows  :  An  atom  is  the  smallest  particle  of  an  element 
which  can  exist.  A  molecule  is  the  smallest  particle  of  a 
compound  which  can  exist.  It  will  be  shown  in  a  subsequent 
chapter  that  sometimes  two  or  more  atoms  of  the  same 
element  unite  with  each  other  to  form  molecules  of  the 
element.  While  the  term  atom,  therefore,  is  applicable 
only  to  elements,  the  term  molecule  is  applicable  both  to 
elements  and  compounds. 


THE  ATOMIC  THEORY  63 

The  atomic  hypothesis  and  the  laws  of  matter.  Suppos- 
ing the  atomic  hypothesis  to  be  true,  let  us  now  see  if  it 
is  in  harmony  with  the  laws  of  matter. 

1.  The  atomic  hypothesis  and  the  law  of  conservation  of 
matter.    It  is  evident  that  if  the  atoms  never  change  their 
masses  in  any  change  which  they  undergo,  the  total  quan- 
tity of  matter  can  never  change  and  the  law  of  conserva- 
tion of  matter  must  follow. 

2.  The  atomic  hypothesis  and  the  law  of  definite  composi- 
tion.   According  to  the  third  supposition,  when  iron  com- 
bines with  sulphur  the  union  is  between  definite  numbers 
of  the  two  kinds  of  atoms.    In  the  simplest  case  one  atom 
of  the  one  element  combines  with  one  atom  of  the  other.   If 
the  sulphur  and  the  iron  atoms  never  change  their  respec- 
tive masses  when  they  unite  to  form  a  molecule  of  iron  sul- 
phide, all  iron  sulphide  molecules  will  have  equal  amounts 
of  iron  in  them  .and  also  of  sulphur.    Consequently  any  mass 
made  up  of  iron  sulphide  molecules  will  have  the  same 
fraction  of  iron  by  weight  as  do  the  individual  iron  sul- 
phide  molecules.     Iron   sulphide,   from  whatever  source, 
will  have  the  same  composition,  which  is  in  accordance 
with  the  law  of  definite  composition. 

3.  The  atomic  hypothesis  and  the  laiv  of  multiple  propor- 
tion. But  this  simplest  case  may  not  always  be  the  only  one. 
Under  other  conditions  one  atom  of  iron  might  combine 
with  two  of  sulphur  to  form  a  molecule  of  a  second  com- 
pound.   In  such   a  case  the  one  atom  of  iron  would   be 
in  combination  with  twice  the  mass  of  sulphur  that  is  in 
the  first  compound,  since  the  sulphur  atoms  all  have  equal 
masses.    What  is  true  for  one  molecule  will  be  true  for  any 
number  of  them  ;   consequently  when  such  quantities  of 
these  two  compounds  are  selected  as  are  found  to  contain 


64       AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

the  same  amount  of  iron,  the  one  will  contain  twice  as 
much  sulphur  as  the  other. 

The  combination  between  the  atoms  may  of  course  take 
place  in  other  simple  ratios.  For  example,  two  atoms  of 
one  element  might  combine  with  three  or  with  five  of  the 
other.  In  all  such  cases  it  is  clear  that  the. law  of  multiple 
proportion  must  hold  true.  For  on  selecting  such  numbers 
of  the  two  kinds  of  molecules  as  have  the  same  number  of 
the  one  kind  of  atoms,  the  numbers  of  the  other  kind  of 
atoms  will  stand  in  some  simple  ratio  to  each  other,  and 
their  weights  will  therefore  stand  in  the  same  simple  ratio. 

Testing  the  hypothesis.  Efforts  have  been  made  to  find 
compounds  which  do  not  conform  to  these  laws,  but  all 
such  attempts  have  resulted  in  failure.  If  such  compounds 
should  be  found,  the  laws  would  be  no  longer  true,  and  the 
hypothesis  of  Dalton  would  cease  to  possess  value.  When 
an  hypothesis  has  been  tested  in  every  way  in  which  experi- 
ment can  test  it,  and  is  still  found  to  be  in  harmony  with 
the  facts  in  the  case,  it  is  termed  a  theory.  We  now  speak 
of  the  atomic  theory  rather  than  of  the  atomic  hypothesis. 

Value  of  a  theory.  The  value  of  a  theory  is  twofold. 
It  aids  in  the  clear  understanding  of  the  laws  of  nature 
because  it  gives  an  intelligent  idea  as  to  why  these  laws 
should  be  in  operation. 

A  theory  also  leads  to  discoveries.  It  usually  happens 
that  in  testing  a  theory  much  valuable  work  is  done,  and 
many  new  facts  are  discovered.  Almost  any  theory  in  ex- 
plaining given  laws  will  involve  a  number  of  consequences 
apart  from  the  laws  it  seeks  to  explain.  Experiment  will 
soon  show  whether  these  facts  are  as  the  theory  predicts 
they  will  be.  Thus  Dalton's  atomic  theory  predicted  many 
properties  of  gases  which  experiment  has  since  verified. 


THE  ATOMIC  THEORY  65 

Atomic  weights.  It  would  be  of  great  advantage  in  the 
study  of  chemistry  if  we  could  determine  the  weights  of 
the  different  kinds  of  atoms.  It  is  evident  that  this  can- 
not be  done  directly.  They  are  so  small  that  they  cannot 
be  seen  even  with  a  most  powerful  microscope.  It  is  cal- 
culated that  it  would  take  200,000,000  hydrogen  atoms 
placed  side  by  side  to  make  a  row  one  centimeter  long.  No 
balance  can  weigh  such  minute  objects.  It  is  possible,  how- 
ever, to  determine  their  relative  weights,  —  that  is,  how 
much  heavier  one  is  than  another.  These  relative  weights 
of  the  atoms  are  spoken  of  as  tJie  atomic  weights  of  the 
elements. 

If  elements  were  able  to  combine  in  only  one  way,  —  one 
atom  of  one  with  one  atom  of  another,  —  the  problem  of  de- 
termining the  atomic  weights  would  be  very  simple.  We 
should  merely  have  to  take  some  one  convenient  element 
as  a  standard,  and  find  by  experiment  how  much  of  each 
other  element  would  combine  with  a  fixed  weight  of  it. 
The  ratios  thus  found  would  be  the  same  ratios  as  those 
between  the  atoms  of  the  elements,  and  thus  we  should 
have  their  relative  atomic  weights.  The  law  of  multiple 
proportion  calls  attention  to  the  fact  that  the  atoms  com- 
bine in  other  ratios  than  I  :  I,  and  there  is  no  direct  way 
of  telling  which  one,  if  any,  of  the  several  compounds  in  a 
given  case  is  the  one  consisting  of  a  single  atom  of  each 
element. 

If  some  way  were  to  be  found  of  telling  how  much 
heavier  the  entire  molecule  of  a  compound  is  than  the  atom 
chosen  as  a  standard,  —  that  is,  of  determining  the  molec- 
ular weights  of  compounds, — the  problem  could  be  solved, 
though  its  solution  would  not  be  an  entirely  simple  matter. 
There  are  ways  of  determining  the  molecular  weights  of 


66       AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

compounds,  and  there  are  other  experiments  which  throw 
light  directly  upon  the  relative  weights  of  the  atoms.  These 
methods  cannot  be  described  until  the  facts  upon  which 
they  rest  have  been  studied.  It  will  be  sufficient  for  the 
present  to  assume  that  these  methods  are  trustworthy. 

Standard  for  atomic  weights.  Since  the  atomic  weights 
are  merely  relative  to  some  one  element  chosen  as  a  stand- 
ard, it  is  evident  that  any  one  of  the  elements  may  serve 
as  this  standard  and  that  any  convenient  value  may  be 
assigned  to  its  atom.  At  one  time  oxygen  was  taken  as 
this  standard,  with  the  value  100,  and  the  atomic  weights 
of  the  other  elements  were  expressed  in  terms  of  this  stand- 
ard. It  would  seem  more  rational  to  take  the  element  of 
smallest  atomic  weight  as  the  standard  and  give  it  unit 
value ;  accordingly  hydrogen  was  taken  as  the  standard 
with  an  atomic  weight  of  I.  Very  recently,  however,  this 
unit  has  been  replaced  by  oxygen,  with  an  atomic  weight 
of  1 6. 

Why  oxygen  is  chosen  as  the  standard  for  atomic  weights. 
In  the  determination  of  the  atomic  weight  of  an  element  it 
is  necessary  to  find  the  weight  of  the  element  which  com- 
bines with  a  definite  weight  of  another  element,  preferably 
the  element  chosen  as  the  standard.  Since  oxygen  com- 
bines with  the  elements  far  more  readily  than  does  hydro- 
gen to  form  definite  compounds,  it  is  far  better  adapted  for 
the  standard  element,  and  has  accordingly  replaced  hydro- 
gen as  the  standard.  Any  definite  value  might  be  given  to 
the  weight  of  the  oxygen  atom.  In  assigning  a  value  to  it, 
however,  it  is  convenient  to  choose  a  whole  number,  and 
as  small  a  number  as  possible  without  making  the  atomic 
weight  of  any  other  element  less  than  unity.  For  these 
reasons  the  number  16  has  been  chosen  as  the  atomic 


THE  ATOMIC  THEORY  67 

weight  of  oxygen.  This  makes  the  atomic  weight  of  hydro- 
gen equal  to  1.008,  so  that  there  is  but  little  difference 
between  taking  oxygen  as  16  and  hydrogen  as  I  for  the 
unit. 

The  atomic  weights  of  the  elements  are  given  in  the 
Appendix. 

EXERCISES 

1.  Two  compounds  were  found  to  have  the  following  compositions  : 
(a)  oxygen  =  69.53%,  nitrogen  =  30.47%;  (£)  oxygen  =  53.27%,  nitro- 
gen =  46.73%.     Show  that  the  law  of  multiple  proportion  holds  in 
this  case. 

2.  Two  compounds  were  found  to  have  the  following  compositions  : 
(a)   oxygen  =  43.64%,  phosphorus  =  56.36%  ;    (£)   oxygen  =  56.35%, 
phosphorus  =  43-65%.    Show  that  the  law  of  multiple  proportion  holds 
in  this  case. 

3.  Why  did  Dalton  assume  that  all  the  atoms  of  a  given  element 
have  the  same  weight? 


CHAPTER  VI 
CHEMICAL  EQUATIONS  AND  CALCULATIONS 

Formulas.  Since  the  molecule  of  any  chemical  com- 
pound consists  of  a  definite  number  of  atoms,  and  this  num- 
ber never  changes  without  destroying  the  identity  of  the 
compound,  it  is  very  convenient  to  represent  the  composi- 
tion of  a  compound  by  indicating  the  composition  of  its 
molecules.  This  can  be  done  very  easily  by  using  the  sym- 
bols of  the  atoms  to  indicate  the  number  and  the  kind  of 
the  atoms  which  constitute  the  molecule.  HgO  will  in  this 
way  represent  mercuric  oxide,  a  molecule  of  which  has 
been  found  to  contain  'i  atom  each  of  mercury  and  oxy- 
gen. H2O  will  represent  water,  the  molecules  of  which 
consist  of  i  atom  of  oxygen  and  2  of  hydrogen,  the  sub- 
script figure  indicating  the  number  of  the  atoms  of  the  ele- 
ment whose  symbol  precedes  it.  H2SO4  will  stand  for 
sulphuric  acid,  the  molecules  of  which  contain  2  atoms  of 
hydrogen,  i  of  sulphur,  and  4  of  oxygen.  The  combination 
of  symbols  which  represents  the  molecule  of  a  substance 
is  called  \\&  formula. 

Equations.  When  a  given  substance  undergoes  a  chemi- 
cal change  it  is  possible  to  represent  this  change  by  the 
use  of  such  symbols  and  formulas.  In  a  former  chapter  it 
was  shown  that  mercuric  oxide  decomposes  when  heated 
to  form  mercury  and  oxygen.  This  may  be  expressed  very 
briefly  in  the  form  of  the  equation 


(i) 

68 


CHEMICAL  EQUATIONS  AND  CALCULATIONS      69 

When  water  is  electrolyzed  two  new  substances,  hydro- 
gen and  oxygen,  are  formed  from  it.  This  statement  in 
the  form  of  an  equation  is 

(2)    H2O  =  2H+O. 

The  coefficient  before  the  symbol  for  hydrogen  indicates 
that  a  single  molecule  of  water  yields  two  atoms  of  hydro- 
gen on  decomposition. 

In  like  manner  the  combination  of  sulphur  with  iron  is 
expressed  by  the  equation 

(3)    Fe  +  S  =  FeS. 

The  decomposition  of  potassium  chlorate  by  heat  takes 
place  as  represented  by  the  equation 

(4)    KC1O3  -  KC1  +  3  O. 

Reading  of  equations.  Since  equations  are  simply  a  kind 
of  shorthand  way  of  indicating  chemical  changes  which 
occur  under  certain  conditions,  in  reading  an  equation  the 
full  statement  for  which  it  stands  should  be  given.  Equa- 
tion (i)  should  be  read,  "Mercuric  oxide  when  heated 
gives  mercury  and  oxygen";  equation  (2)  is  equivalent  to 
the  statement,  "  When  electrolyzed,  water  produces  hydro- 
gen and  oxygen"  ;  equation  (3),  "  When  heated  together 
iron  and  sulphur  unite  to  form  iron  sulphide  ";  equation  (4), 
"  Potassium  chlorate  when  heated  yields  potassium  chloride 
and  oxygen." 

Knowledge  required  for  writing  equations.  In  order  to 
write  such  equations  correctly,  a  considerable  amount  of 
exact  knowledge  is  required.  Thus,  in  equation  (i)  the  fact 
that  red  oxide  of  mercury  has  the  composition  represented 
by  the  formula  HgO,  that  it  is  decomposed  by  heat,  that 
in  this  decomposition  mercury  and  oxygen  are  formed  and 


70       AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

no  other  products,  —  all  these  facts  must  be  ascertained  by 
exact  experiment  before  the  equation  can  be  written.  An 
equation  expressing  these  facts  will  then  have  much  value. 
Having  obtained  an  equation  describing  the  conduct  of 
mercuric  oxide  on  being  heated,  it  will  not  do  to  assume 
that  other  oxides  will  behave  in  like  manner.  Iron  oxide 
(FeO)  resembles  mercuric  oxide  in  many  respects,  but  it 
undergoes  no  change  at  all  when  heated.  Manganese 
dioxide,  the  black  substance  used  in  the  preparation  of 
oxygen,  has  the  formula  MnO2.  When  this  substance  is 
heated  oxygen  is  set  free,  but  the  metal  manganese  is  not 
liberated  ;  instead,  a  different  oxide  of  manganese  contain- 
ing less  oxygen  is  produced.  The  equation  representing 

the  reaction  is 

3  MnO2  =  Mn3O4  +  2  O. 

Classes  of  reactions.  When  a  chemical  change  takes 
place  in  a  substance  the  substance  is  said  to  undergo  a 
reaction.  Although  a  great  many  different  reactions  will 
be  met  in  the  study  of  chemistry,  they  may  all  be  grouped 
under  the  following  heads. 

1.  Addition.     This    is    the   simplest    kind    of   chemical 
action.    It  consists  in  the  union  of  two  or  more  substances 
to  produce  a  new  substance.    The  combination  of  iron  with 
sulphur  is  an  example  : 

Fe  +  S  =  FeS. 

2.  Decomposition.    This  is  the  reverse  of  addition,  the 
substance  undergoing  reaction  being  parted  into  its  con- 
stituents.   The    decomposition    of    mercuric    oxide    is    an 
example:  HgO  =  Hg  +  O. 

3.  Substitution.    It  is  sometimes  possible  for  an  element 
in  the  free  state  to  act  upon  a  compound  in  such  a  way  that 


CHEMICAL  EQUATIONS  AND  CALCULATIONS      71 

it  takes  the  place  of  one  of  the  elements  of  the  compound, 
liberating  it  in  turn.  In  the  study  of  the  element  hydrogen 
it  was  pointed  out  that  hydrogen  is  most  conveniently  pre- 
pared by  the  action  of  sulphuric  or  hydrochloric  acid  upon 
zinc.  When  sulphuric  acid  is  used  a  substance  called  zinc 
sulphate,  having  the  composition  represented  by  the  formula 
ZnSO4,  is  formed  together  with  hydrogen.  The  equation  is 
Zn  +  H2SO4  =  ZnSO4  +  2  H. 

When  hydrochloric  acid  is  used  zinc  chloride  and  hydro- 
gen are  the  products  of  reaction  : 

Zn  +  2  HC1  =  ZnCl2  +  2  H. 

When  iron  is  used  in  place  of  zinc  the  equation  is 
Fe  +  H2SO4  =  FeSO4  +  2  H. 

These  reactions  are  quite  similar,  as  is  apparent  from  an 
examination  of  the  equations.  In  each  case  i  atom  of  the 
metal  replaces  2  atoms  of  hydrogen  in  the  acid,  and  the 
hydrogen  escapes  as  a  gas.  When  an  element  in  the  free 
state,  such  as  the  zinc  in  the  equations  just  given,  takes  the 
place  of  some  one  element  in  a  compound,  setting  it  free 
from  chemical  combination,  the  act  is  called  substitution. 

Other  reactions  illustrating  substitution  are  the  action 
of  sodium  on  water, 

Na  +  H20  =  NaOH  +  H  ; 
and  the  action  of  heated  iron  upon  water, 

3  Fe  +  4  H2O  =  Fe3O4  +  8  H. 

4.  Double  decomposition.  When  barium  dioxide  (BaO2) 
is  treated  with  sulphuric  acid  two  compounds  are  formed, 
namely,  hydrogen  dioxide  (H2O2)  and  barium  sulphate 
(BaSO4).  The  equation  is 

BaO2  +  H2SO4  =  BaSO4  +  H2O2. 


72        AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

In  this  reaction  it  will  be  seen  that  the  two  elements 
barium  and  hydrogen  simply  exchange  places.  Such  a^ 
reaction  is  called  a  double  decomposition.  We  shall  meet 
with  many  examples  of  this  kind  of  chemical  reactions. 

Chemical  equations  are  quantitative.  The  use  of  sym- 
bols and  formulas  in  expressing  chemical  changes  has 
another  great  advantage.  Thus,  according  to  the  equation 

H2O  =  2  H  +  O, 

I  molecule  of  water  is  decomposed  into  2  atoms  of  hydro- 
gen and  i  atom  of  oxygen.  But,  as  we  have  seen,  the  rela- 
tive weights  of  the  atoms  are  known,  that  of  hydrogen  being 
1.008,  while  that  of  oxygen  is  16.  The  molecule  of  water, 
being  composed  of  2  atoms  of  hydrogen  and  I  atom  of  oxy- 
gen, must  therefore  weigh  relatively  2.016  +  16,  or  18.01 6. 
The  amount  of  hydrogen  in  this  molecule  must  be  ^—.,  or 
1 1 . 1 8  %  of  the  whole,  while  the  amount  of  oxygen  must  be 
— -,  or  88.82%  of  the  whole.  Now,  since  any  definite  quan- 
tity of  water  is  simply  the  sum  of  a  great  many  molecules 
of  water,  it  is  plain  that  the  fractions  representing  the 
relative  amounts  of  hydrogen  and  oxygen  present  in  a 
molecule  must  likewise  ekpress  the  relative  amounts  of 
hydrogen  and  oxygen  present  in  any  quantity  of  water. 
Thus,  for  example,  in  20  g.  of  water  there  are  ^j,  x  20, 
or  2.238  g.  of  hydrogen,  and  -^~  x  20,  or  17.762  g.  of 
oxygen.  These  results  in  reference  to  the  composition  of 
water  of  course  agree  exactly  with  the  facts  obtained  by 
the  experiments  described  in  the  chapter  on  water,  for  it  is 
because  of  those  experiments  that  the  values  i  .008  and  16 
are  given  to  hydrogen  and  oxygen  respectively. 

It  is  often  easier  to  make  calculations  of  this  kind  in  the 
form  of  a  proportion  rather  than  by  fractions.  Since  the 


CHEMICAL  EQUATIONS  AND  CALCULATIONS      73 

molecule  of  water  and  the  two  atoms  of  hydrogen  which  it 
contains  have  the  ratio  by  weight  of  18.016:2.016,  any 
mass  of  water  has  the  same  ratio  between  its  total  weight 
and  the  weight  of  the  hydrogen  in  it.  Hence,  to  find  the 
number  of  grams  (x)  of  hydrogen  in  20  g.  of  water,  we 
have  the  proportion 

1 8.016  :  2.016  ::  20  g.  \x  (grams  of  hydrogen). 
Solving  for  x,  we  get  2.238  for  the  number  of  grams  of 
hydrogen.    Similarly,  to  find  the  amount  (x)  of  oxygen  pres- 
ent in  the  20  g.  of  water,  we  have  the  proportion 

18.016  :  16  ::  20  \x\ 
from  which  we  find  that  x  =  17.762  g. 

Again,  suppose  we  wish  to  find  what  weight  of  oxygen  can  be 
obtained  from  15  g.  of  mercuric  oxide.  The  equation  representing 
the  decomposition  of  mercuric  oxide  is 

HgO  =  Hg  +  O. 

The  relative  weights  of  the  mercury  and  oxygen  atoms  are  respec- 
tively 200  and  16.  The  relative  weight  of  the  mercuric  oxide  mole- 
cule must  therefore  be  the  sum  of  these,  'or  216.  The  molecule  of 
mercuric  oxide  and  the  atom  of  oxygen  which  it  contains  have  the 
ratio  216  :i6.  This  same  ratio  must  therefore  hold  between  the 
weight  of  any  given  quantity  of  mercuric  oxide  and  that  of  the  oxygen 
which  it  contains.  Hence,  to  find  the  weight  of  oxygen  in  15  g.  of 
mercuric  oxide,  we  have  the  proportion 

216  :  16  : :  15  :  x  (grams  of  oxygfen). 

On  the  other  hand,  suppose  we  wish  to  prepare,  say,  20  g.  of 
oxygen.  The  problem  is  to  find  out  what  weight  of  mercuric  oxide 
will  yield  20  g.  of  oxygen.  The  following  proportion  evidently  holds 

216  :  1 6  : :  x  (grams  of  mercuric  oxide)  :  20 ; 

from  which  we  get  ;r=  270. 

In  the  preparation  of  hydrogen  by  the  action  of  sulphuric  acid 
upon  zinc,  according  to  the  equation, 

Zn  +  H2SO4  =  ZnSO4  +  2  H, 


74        AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

suppose  that  50  g.  of  zinc  are  available ;  let  it  be  required  to  cal- 
culate the  weight  of  hydrogen  which  can  be  obtained.  It  will  be 
seen  that  i  atom  of  zinc  will  liberate  2  atoms  of  hydrogen.  The 
ratio  by  weight  of  a  zinc  to  an  hydrogen  atom  is  65.4  :  1.008  ;  of  I 
zinc  atom  to  2  hydrogen  atoms,  65.4:2.016.  Zinc  and  hydrogen 
will  be  related  in  this  reaction  in  this  same  ratio,  however  many 
atoms  of  zinc  are  concerned.  Consequently  in  the  proportion 

65.4  :  2.016  :  :  50  :  x, 

x  will  be  the  weight  of  hydrogen  set  free  by  50  g.  of  zinc.  The 
weight  of  zinc  sulphate  produced  at  the  same  time  can  be  found 
from  the  proportion 

65.4  :  161146  : :  50  :  jr, 

where  161.46  is  the  molecular  weight  of  the  zinc  sulphate,  and  x  the 
weight  of  zinc  sulphate  formed.  In  like  manner,  the  weight  of  sul- 
phuric acid  used  up  can  be  calculated  from  the  proportion 

65.4:98.076  : :  50  :  x. 

These  simple  calculations  are  possible  because  the  sym- 
bols and  formulas  in  the  equations  represent  the  relative 
weights  of  the  substances  concerned  in  a  chemical  reaction. 
When  once  the  relative  weights  of  the  atoms  have  been 
determined,  and  it  has  been  agreed  to  allow  the  symbols  to 
stand  for  these  relative  weights,  an  equation  or  formula 
making  use  of  the  symbols  becomes  a  statement  of  a  defi- 
nite numerical  fact,  and  calculations  can  be  based  on  it. 

Chemical  equations  not  algebraic.  Although  chemical 
equations  are  quantitative,  it  must  be  clearly  understood 
that  they  are  not  algebraic.  A  glance  at  the*  equations 

7  +  4=11,  8  +  5-9  +  4 

will  show  at  once  that  they  are  true.    The  equations 
HgO  =  Hg  +  O,     FeO  =  Fe  +  O 

are  equally  true  in  an  algebraic  sense,  but  experiment  shows 
that  only  the  first  is  true  chemically,  for  iron  oxide  (FeO) 


CHEMICAL  EQUATIONS  AND  CALCULATIONS      75 

cannot  be  directly  decomposed  into  iron  and  oxygen.  Only 
such  equations  as  have  been  found  by  careful  experiment 
to  express  a  real  chemical  transformation,  true  both  for  the 
kinds  of  substances  as  well  as  for  the  weights,  have  any 
value. 

Chemical  formulas  and  equations,  therefore,  are  a  concise 
way  of  representing  qualitatively  and  quantitatively  facts 
which  have  been  found  by  experiment  to  be  true  in  reference 
to  the  composition  of  substances  and  the  changes  which  they 
undergo, 

Formulas  representing  water  of  crystallization.  An  ex- 
amination of  substances  containing  water  of  crystallization 
has  shown  that  in  every  case  the  water  is  present  in  such 
proportion  by  weight  as  can  readily  be  represented  by  a  for- 
mula. For  example,  copper  sulphate  (CuSO4)  and  water 
combine  in  the  ratio  of  I  molecule  of  the  sulphate  to  5  of 
water  ;  calcium  sulphate  (CaSO4)  and  water  combine  in  the 
ratio  i  :  2  to  form  gypsum.  These  facts  are  expressed  by 
writing  the  formulas  for  the  two  substances  with  a  period 
between  them.  Thus  the  formula  for  crystallized  copper 
sulphate  is  CuSO4-  5  H2O  ;  that  of  gypsum  is  CaSO4-  2  H2O. 

Heat  of  reaction.  Attention  has  frequently  been  directed 
to  the  fact  that  chemical  changes  are  usually  accompanied 
by  heat  changes.  In  general  it  has  been  found  that  in  every 
chemical  action  heat  is  either  absorbed  or  given  off.  By 
adopting  a  suitable  unit  for  the  measurement  of  heat,  the 
heat  change  during  a  chemical  reaction  can  be  expressed  in 
the  equation  for  the  reaction. 

Heat  cannot  be  measured  by  the  use  of  a  thermometer 
alone,  since  the  thermometer  measures  the  intensity  of  heat, 
not  its  quantity.  The  easiest  way  to  measure  a  quantity  of 
heat  is  to  note  how  warm  it  will  make  a  definite  amount  of 


76       AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

a  given  substance  chosen  as  a  standard.  Water  has  been 
chosen  as  the  standard,  and  the  unit  of  heat  is  called  a 
calorie.  A  calorie  is  defined  as  the  amount  of  heat  required 
to  raise  the  temperature  of  one  gram  of  water  one  degree. 

By  means  of  this  unit  it  is  easy  to  indicate  the  heat 
changes  in  a  given  chemical  reaction.  The  equation 

2  H  +  O  =  H2O  +  68,300  cal. 

means  that  when  2.016  g.  of  hydrogen  combine  with  16  g. 
of  oxygen,  18.01 6  g.  of  water  are  formed  and  68,300  cal. 
are  set  free. 

C  4-  2  S  =  CS2  —  19,000  cal. 

means  that  an  expenditure  of  19,000  cal.  is  required  to 
cause  12  g.  of  carbon  to  unite  with  64.12  g.  of  sulphur  to 
form  76.12  g.  of  carbon  disulphide.  In  these  equations  it 
will  be  noted  that  the  symbols  stand  for  as  many  grams 
of  the  substance  as  there  are  units  in  the  weights  of  the 
atoms  represented  by  the  symbols.  This  is  always  under- 
stood to  be  the  case  in  equations  where  the  heat  of  reaction 
is  given. 

Conditions  of  a  chemical  action  are  not  indicated  by  equa- 
tions. Equations  do  not  tell  the  conditions  under  which  a 
reaction  will  take  place.  The  equation 

HgO  =  Hg  +  O 

does  not  tell  us  that  it  is  necessary  to  keep  the  mercuric 
oxide  at  a  high  temperature  in  order  that  the  decomposition 
may  go  on.  The  equation 

Zn  +  2  HC1  -  ZnCl2  +  2  H 

in  no  way  indicates  the  fact  that  the  hydrochloric  acid  must 
be  dissolved  in  water  before  it  will  act  upon  the  zinc.  From 
the  equation  H  +  Cl  =  HC1 


CHEMICAL  EQUATIONS  AND  CALCULATIONS    77 

it  would  not  be  suspected  that  the  two  gases  hydrogen  and 
chlorine  will  unite  instantly  in  the  sunlight,  but  will  stand 
mixed  in  the  dark  a  long  time  without  change.  It  will 
therefore  be  necessary  to  pay  much  attention  to  the  details 
of  the  conditions  under  which  a  given  reaction  occurs,  as 
well  as  to  the  expression  of  the  reaction  in  the  form  of 
an  equation. 

EXERCISES 

1.  Calculate  the  percentage  composition  of  the  following  sub- 
stances :   (a)   mercuric  oxide  ;    (£)   potassium  chlorate ;  (c}   hydro- 
chloric acid;    (</)  sulphuric  acid.      Compare    the  results   obtained 
with  the  compositions  as  given  in  Chapters  II  and  III. 

2.  Determine   the  percentage  of  copper,   sulphur,  oxygen,   and 
water  in  copper  sulphate  crystals.    What  weight  of  water  can  be 
obtained  from  150  g.  of  this  substance? 

3.  What  weight  of  zinc  can  be  dissolved  in  10  g.  of  sulphuric 
acid  ?    How  much  zinc  sulphate  will  be  formed  ? 

4.  How  many  liters  of  hydrogen  measured  under  standard  condi- 
tions can  be  obtained  from   the  action  of  8  g.  of  iron  on  10  g.  of 
sulphuric  acid?    How  much  iron  sulphate  (FeSO4)  will  be  formed? 

5.  log.  of  zinc  were  used  in  the  preparation  of  hydrogen  ;  what 
weight  of  iron  will  be  required  to  prepare  an  equal  volume  ? 

6.  How  many  grams  of  barium  dioxide  will  be  required  to  prepare 
i  kg.  of  common  hydrogen  dioxide  solution  ?    What  weight  of  barium 
sulphate  will  be  formed  at  the  same  time  ? 

7.  What   weight  of   the  compound   Mn3O4  will   be  formed  by 
strongly  heating  25  g.  of   manganese   dioxide  ?      What  volume   of 
oxygen  will  be  given  off  at  the  same  time,  measured  under  standard 
conditions  ? 

8.  (a)  What  is  the  weight  of  100  1.  of  hydrogen  measured  in  a 
laboratory  in  which  the  temperature  is  20°  and  pressure  750  mm.  ? 
(£)  What  weight   of   sulphuric   acid   is   necessary  to   prepare   this 
amount  of  hydrogen?    (V)  The  density  of  sulphuric  acid  is  1.84. 
Express  the  acid  required  in  (6)  in  cubic  centimeters. 

9.  What  weight  of  potassium  chlorate  is  necessary  to  furnish 
sufficient  oxygen    to   fill   four    200  cc.  bottles   in  your    laboratory 
(the  gas  to  be  collected  over  water)  ? 


CHAPTER  VII 

NITROGEN    AND    THE    RARE    ELEMENTS:    ARGON, 

HELIUM,  NEON,  KRYPTON,  XENON  /^ 

— -^ 

Historical.  Nitrogen  was  discovered  by  the  English  chem- 
ist Rutherford  in  1772.  A  little  later  Scheele  showed  it 
to  be  a  constituent  of  air,  and  Lavoisier  gave  it  the  name 
azote,  signifying  that  it  would  not  support  life.  The  name 
nitrogen  was  afterwards  given  it  because  of  its  presence  in 
saltpeter  or  niter.  The  term  azote  and  symbol  Az  are  still 
retained  by  the  French  chemists. 

Occurrence.  Air  is  composed  principally  of  oxygen  and 
nitrogen  in  the  free  state,  about  78  parts  by  volume  out  of 
every  100  parts  being  nitrogen.  Nitrogen  also  occurs  in 
nature  in  the  form  of  potassium  nitrate  (KNO3)  —  com- 
monly called  saltpeter  or  niter  —  as  well  as  in  sodium  nitrate 
(NaNO3).  Nitrogen  is  also  an  essential  constituent  of  all 
living  organisms  ;  for  example,  the  human  body 
about  2.4%  of  nitrogen. 

Preparation  from  air.  Nitrogen  can  be  prepared  from  air 
by  the  action  of  some  substance  which  will  combine  with 
the  oxygen,  leaving  the  nitrogen  free.  Such  a  substance 
must  be  chosen,  however,  as  will  combine  with  the  oxygen 
to  form  a  product  which  is  not  a  gas,  and  which  can  be 
readily  separated  from  the  nitrogen.  The  substances  most 
commonly  used  for  this  purpose  are  phosphorus  and  copper. 

i.  By  the  action  of  phosphorus.  The  method  used  for 
the  preparation  of  nitrogen  by  the  action  of  phosphorus  is 
as  follows : 

78 


(C       *L  "  ei*u>  iii  O*«  j 
V     ^^      -Hf  \ 

-- NfrlfSSqEN  AND  THE  RARE  ELEMENTS 


79 


A 


- 

^*= 

i 

—_ 

=-_ 

-, 

_-=: 

- 

5 

^ 

~~— 

r£ 

FIG. 26 


The  phosphorus  is  placed  in  a  little  porcelain  dish  supported  on 
a  cork  and  floated  on  water  (Fig.  26).  It  is  then  ignited  by  contact 
with  a  hot  wire,  and  immediately  a  bell  jar  or  bottle  is  brought  over 
it  so  as  to  confine  a  portion  of  the  air.  The  phosphorus  combines 
with  the  oxygen  to  form  an  oxide  of  phosphorus,  known  as  phos- 
phorus pentoxide.  This  is  a  white  solid  which  floats  about  in  the 
bell  jar,  but  in  a  short  time  it  is  all 
absorbed  by  the  water,  leaving  the 
nitrogen.  The  withdrawal  of  the 
oxygen  is  indicated  by  the  rising 
of  the  water  in  the  bell  jar. 

2.  By  the  action  of  copper. 
The  oxygen  present  in  the  air 
may  also  be  removed  by  pass- 
ing air  slowly  through  a 

heated  tube  containing  copper.  The  copper  combines  with 
the  oxygen  to  form  copper  oxide,  which  is  a  solid.  The 
nitrogen  passes  on  and  may  be  collected  over  water. 

Nitrogen  obtained  from  air  is  not  pure.  Inasmuch  as  air,  in  addi- 
tion to  oxygen  and  nitrogen,  contains  small  amounts  of  other  gases, 
and  since  the  phosphorus  as  well  as  the  copper  removes,  only  the 
oxygen,  it  is  evident  that  the  nitrogen  obtained  by  these  methods  is 
never  ^ftte  pure.  About  i%  of  the  product  is  composed  of  other 
gases,  from  which  it  is  very  difficult  to  separate  the  nitrogen.  The- 
impure  nitrogen  so  obtained  may,  however,  be  used  for  a  study  of 
most  of  the  properties  of  nitrogen,  since  these  are  not  materially 
affected  by  the  presence  of  the  other  gases. 

Preparation  from  compounds  of  nitrogen.  Pure  nitrogen 
may  be  obtained  from  certain  compounds  of  the  element. 
Thus,  if  heat  is  applied  to  the  compound  ammonium  nitrite 
(NH4NO2),  the  change  represented  in  the  following  equa- 
tion takes  place  : 

NH4NO2  =  2  H2O  +  2  N, 


80       AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Physical  properties.  Nitrogen  is  similar  to  oxygen  and 
hydrogen  in  that  it  is  a  colorless,  odorless,  and  tasteless  gas. 
One  liter  of  nitrogen  weighs  1.2501  g.  It  is  almost 
insoluble  in  water.  It  can  be  obtained  in  the  form  of  a 
colorless  liquid  having  a  boiling  point  of —  195°  at  ordinary 
pressure.  At  — 214°  it  solidifies. 

Chemical  properties.  Nitrogen  is  characterized  by  its  in- 
ertness. It  is  neither  combustible  nor  a  supporter  of  com- 
bustion. At  ordinary  temperatures  it  will  not  combine 
directly  with  any  of  the  elements  except  under  rare  con- 
ditions. At  higher  temperatures  it  combines  with  mag- 
nesium, lithium,  titanium,  and  a  number  of  other  elements. 
The  compounds  formed  are  called  nitrides,  just  as  com- 
pounds of  an  element  with  oxygen  are  called  oxides.  When 
it  is  mixed  with  oxygen  and  subjected  to  the  action  of  elec- 
tric sparks,  the  two  'gases  slowly  combine  forming  oxides  of 
nitrogen.  A  mixture  of  nitrogen  and  hydrogen  when  treated 
similarly  forms  ammonia,  a  gaseous  compound  of  nitrogen 
and  hydrogen.  Since  we  are  constantly  inhaling  nitrogen, 
it  is  evident  that  it  is  not  poisonous.  Nevertheless  life 
would  be  impossible  in  an  atmosphere  of  pure  nitrogen  on 
account  of  the  exclusion  of  the  necessary  oxygen. 

Argon,  helium,  neon,  krypton,  xenon.  These  are  all  rare  elements 
occurring  in  the  air  in  very  small  quantities.  Argon,  discovered  in 
1894,  was  the  first  one  obtained.  Lord  Rayleigh,  an  'English  sci- 
entist, while  engaged  in  determining  the  exact  weights  of  various 
gases,  observed  that  the  nitrogen  obtained  from  the  air  is  slightly 
heavier  than  pure  nitrogen  obtained  from  its  compounds.  After 
repeating  his  experiments  many  times,  always  with  the  same  results, 
Rayleigh  finally  concluded  that  the  nitrogen  which  he  had  obtained 
from  the  air  was  not  pure,  but  was  mixed  with  a  small  amount  of 
some  unknown  gas,  the  density  of  which  is  greater  than  that  of 
nitrogen.  Acting  on  this  assumption,  Rayleigh,  together  with  the 


NITROGEN  AND  THE  RARE  ELEMENTS         8 1 

English  chemist  Ramsay,  attempted  to  separate  the  nitrogen  from 
the  unknown  gas.  Knowing  that  nitrogen  would  combine  with  mag- 
nesium, they  passed  the  nitrogen  obtained  from  the  air  and  freed 
from  all  known  substances  through  tubes  containing  magnesium 
heated  to  the  necessary  temperature.  After  repeating  this  operation, 
they  finally  succeeded  in  obtaining  from  the  atmospheric  nitrogen 
a  small  volume  of  gas  which  would  not  combine  with  magnesium 
and  hence  could  not  be  nitrogen.  This  proved  to  be  a  new  element, 
to  which  they  gave  the  name  argon.  As  predicted,  this  new  element 
was  found  to  be  heavier  than  nitrogen,  its  density  as  compared  with 
hydrogen  as  a  standard  being  approximately  20,  that  of  nitrogen 
being  only  14.  About  i%  of  the  atmospheric  nitrogen  proved  to  be 
argon.  The  new  element  is  characterized  by  having  no  affinity  for 
other  elements.  Even  under  the  most  favorable  conditions  it  has 
not  been  made  to  combine  with  any  other  element.  On  this  account 
it  was  given  the  name  argon,  signifying  lazy  or  idle.  Like  nitro- 
gen, it  is  colorless,  odorless,  and  tasteless.  It  has  been  liquefied 
and  solidified.  Its  boiling  point  is  — 187°. 

Helium  was  first  found  in  the  gases  expelled  from  certain  minerals 
by  heating.  Through  the  agency  of  the  spectroscope  it  had  been 
known  to  exist  in  the  sun  long  before  its  presence  on  the  earth  had 
been  demonstrated,  —  a  fact  suggested  by  the  name  helium,  signify- 
ing the  sun.  Its  existence  in  traces  in  the  atmosphere  has  also  been 
proven.  It  is  the  only  gas  which  has  not  yet  been  liquefied,  so  that 
its  boiling  point  must  be  below  that  of  hydrogen. 

The  remaining  elements  of  this  group  —  neon,  krypton,  and  xenon 
—  have  been  obtained  from  liquid  air.  When  liquid  air  is  allowed 
to  boil,  the  constituents  which  are  the  most  difficult  to  liquefy,  and 
which  therefore  have  the  lowest  boiling  points,  vaporize  first,  fol- 
lowed by  the  others  in  the  order  of  their  boiling  points.  It  is  possible 
in  this  way  to  make  at  least  a  partial  separation  of  the  air  into  its  con- 
stituents, and  Ramsay  thus  succeeded  in  obtaining  from  liquid  air  not 
only  the  known  constituents,  including  argon  and  helium,  but  also  the 
new  elements,  neon,  krypton,  and  xenon.  These  elements,  as  well  as 
helium,  all  proved  to  be  similar  to  argon  in  that  they  are  without 
chemical  activity,  apparently  forming  no  compounds  whatever.  The 
percentages  present  in  the  air  are  very  small.  The  names,  neon, 
krypton,  xenon,  signify  respectively,  new,  hidden,  stranger. 


82        AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


EXERCISES 

1.  How  could  you  distinguish  between  oxygen,   hydrogen,  and 
nitrogen  ? 

2.  Calculate  the  relative  weights   of   nitrogen  and  oxygen ;    of 
nitrogen  and  hydrogen. 

3.  In  the  preparation  of  nitrogen  from  the  air,  how  would  hydro- 
gen do  as  a  substance  for  the  removal  of  the  oxygen  ? 

4.  What  weight  of  nitrogen   can   be  obtained  from  10  1.  of  air 
measured  under  the  conditions  of  temperature  and  pressure  which 
prevail  in  your  laboratory  ? 

5.  How  many  grams  of  ammonium  nitrite  are  necessary  in  the 
preparation  of  20  1.  of  nitrogen  measured  over  water  under  the  con- 
ditions of  temperature  and  pressure  which  prevail  in  your  laboratory? 

6.  If  10  1.  of  air,  measured  under  standard  conditions,  is  passed 
over  100  g.  of  hot  copper,  how  much  will  the  copper  gain  in  weight? 


CHAPTER  VIII 
THE  ATMOSPHERE 

Atmosphere  and  air.  The  term  atmosphere  is  applied  to 
the  gaseous  envelope  surrounding  the  earth.  The  term  air 
is  generally  applied  to  a  limited  portion  of  this_  envelope, 
although  the  two  words  are  often  used  interchangeably. 
Many  references  have  already  been  made  to  the  composition 
and  properties  of  the  atmosphere.  These  statements  must 
now  be  collected  and  discussed  somewhat  more  in  detail. 

Air  formerly  regarded  as  an  element.  Like  water,  air 
was  at  first  regarded  as  elementary  in  character.  Near  the 
close  of  the  eighteenth  century  Scheele,  Priestley,  and 
Lavoisier  showed  by  their  experiments  that  it  is  a  mixture 
of  at  least  two  gases,  —  those  which  we  now  call  oxygen  and 
nitrogen.  By  burning  substances  in  an  inclosed  volume  of 
air  and  noting  the  contraction  in  volume  due  to  the  removal 
of  the  oxygen,  they  were  able  to  determine  with  some  accur- 
acy the  relative  volumes  of  oxygen  and  nitrogen  present 
in  the  air. 

The  constituents  of  the  atmosphere.  The  constituents  of 
the  atmosphere  may  be  divided  into  two  general  groups : 
those  which  are  essential  to  life  and  those  which  are  not 
essential. 

I.  Constituents  essential  to  life.  In  addition  to  oxygen 
and  nitrogen  at  least  two  other  substances,  namely,  carbon 
dioxide  and  water  vapor,  must  be  present  in  the  atmosphere 
in  order  that  life  may  exist.  The  former  of  these  is  a 

83 


84       AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

gaseous  compound  of  carbon  and  oxygen  having  the  for- 
mula CO2.  Its  properties  will  be  discussed  in  detail  in  the 
chapter  on  the  compounds  of  carbon.  Its  presence  in  the 
air  may  be  shown  by  causing  the  air  to  bubble  through 
a  solution  of  calcium  hydroxide  (Ca(OH)2),  commonly 
called  lime  water.  The  carbon  dioxide  combines  with 
the  calcium  hydroxide  in  accordance  with  the  following 

equation : 

Ca(OH)2  +  C02  =  CaC03  +  H2O. 

The  resisting  calcium  carbonate  (CaCO3)  is  insoluble  in 
water  and  separates  in  the  form  of  a  white  powder,  which 
causes  the  solution  to  appear  milky. 

The  presence  of  water  vapor  is  readily  shown  by  its  con- 
densation on  cold  objects  as  well  as  by  the  fact  that  a  bit 
of  calcium  chloride  when  exposed  to  the  air  becomes  moist, 
and  may  even  dissolve  in  the  water  absorbed  from  the  air. 

2.  Constituents  not  essential  to  life.  In  addition  to  the 
essential  constituents,  the  air  contains  small  percentages  of 
various  other  gases,  the  presence  of  which  so  far  as  is 
known  is  not  essential  to  life.  This  list  includes  the  rare 
elements,  argon,  helium,  neon,  krypton,  and  xenon  ;  also 
hydrogen,  ammonia,  hydrogen  dioxide,  and  probably  ozone. 
Certain  minute  forms  of  life  (germs)  are  also  present,  the 
decay  of  organic  matter  being  due  to  their  presence. 

Function  of  each  of  the  essential  constituents,  (i)  The  oxygen 
directly  supports  life  through  respiration.  (2)  The  nitrogen,  on 
account  of  its  inactivity,  serves  to  dilute  the  oxygen,  and  while  con- 
trary to  the  older  views,  it  is  possible  that  life  might  continue  to 
exist  in  the  absence  of  the  atmospheric  nitrogen,  yet  the  conditions 
of  life  would  be  entirely  changed.  Moreover,  nitrogen  is  an  essential 
constituent  of  all  animal  and  'plant  life.  It  was  formerly  supposed 
that  neither  animals  nor  plants  could  assimilate  the  free  nitrogen, 
but  it  has  been  shown  recently  that  the  plants  of  aUeast  one  natural 


THE  ATMOSPHERE 


order,  the  Leguminosae,  to  which  belong  the  beans,  peas,  and  clover, 

have  the  power  of  directly  assimilating  the  free  nitrogen  from  the 

atmosphere.    This  is  accomplished 

through  the   agency  of  groups  of 

bacteria,  which  form  colonies  in  little 

tubercles  on  the  roots  of  the  plants. 

These  bacteria  probably  assist. in  the 

absorption  of  nitrogen  by  changing  the 

free  nitrogen   into  compounds  which 

can  be  assimilated  by  the  plant. 

Fig.   27  shows  the   tubercles   on  the 

roots  of  a  variety  of  bean.     (3)  The 


FIG. 27 


presence  of  water  vapor  in  the  air  is  necessary  to  prevent  excessive 
evaporation  from  both  plants  and  animals.  (4)  Carbon  dioxide  is 
an  essential  plant  food. 

The  quantitative  analysis  of  air.  A  number  of  different 
methods  have  been  devised  for  the  determination  of  the 
percentages  of  the  constituents  present  in  the  atmosphere. 
Among  these  are  the  following. 

I .  Determ ination  of  oxygen .  ( I )  The 
oxygen  is  withdrawn  from  a  measured 
volume  of  air  inclosed  in  a  tube,  by 
means  of  phosphorus. 

To  make  the  determination,  a  graduated 
tube  is  filled  with  water  and  inverted  in  a 
vessel  of  water.  Air  is  introduced  into  the 
tube  until  it  is  partially  filled  with  the  gas. 
The  volume  of  the  inclosed  air  is  carefully 
noted  and  reduced  to  standard  conditions.  A 
smalPpiece  of  phosphorus  is  attached  to  a 
wire  and  brought  within  the  tube  as  shown 
in  Fig.  28.  After  a  few  hours  the  oxygen  in 
the  inclosed  air  will  have  combined  with  the 
phosphorus,  the  water  rising  to  take  its  place.  The  phosphorus  is 
removed  and  the  volume  is  again  noted  and  reduced  to  standard 
conditions.  The  contraction  in  the  volume  of  the  air  is  equal  to  the 
volume  of  oxygen  absorbed. 


FIG.  28 


86       AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

(2)  The    oxygen  may  also  be  estimated  by  passing  a 
measured  volume  of  air  through  a  tube  containing  copper 
heated  to  a  high  temperature.    The  oxygen  in  the  air  com- 
bines with  the  copper  to  form  copper  oxide  (CuO).    Hence 
the  increase  in  the  weight  of  the  copper  equals  the  weight 
of  the  oxygen  in  the  volume  of  air  taken. 

(3)  A  more  accurate  method  is  the  following.    A  eudiom- 
eter tube  is  filled  with  mercury  and  inverted  in  a  vessel 
of  the  same  liquid.    A  convenient  amount  of  air  is  then 
introduced  into  the  tube  and  its  volume  accurately  noted. 
There  is  then  introduced  more  than  sufficient  hydrogen  to 
combine  with  the  oxygen  present  in  the  inclosed  air,  and 
the  volume  is  again  accurately  noted.    The  mixture  is  then 
exploded  by  an  electric  spark,  and  the  volume  is  once  more 
taken.    By  subtracting  this  volume  from  the  total  volume 
of  the  air  and  hydrogen  there  is  obtained  the  contraction 
in  volume  due  to  the  union  of  the  oxygen  and  hydrogen. 
The  volume  occupied  by  the  water  formed  by  the  union  of 
the  two  gases  is  so  small  that  it  may  be  disregarded  in  the 
calculation.     Since  oxygen  and  hydrogen  combine  in  the 
ratio   I  :  2  by  volume,  it   is   evident   that   the  contraction 
in  volume  due  to  the  combination  is  equal  to  the  volume 
occupied  by  the  oxygen  in  the  air  contained  in  the  tube, 
plus  twice  this  volume  of  hydrogen.    In  other  words,  one 
third  of  the  total  contraction  is  equal  to  the  volume  occu- 
pied by  the  oxygen   in   the  inclosed  air.    The    following 
example  will  make  this  clear  : 

Volume  of  air  in  tube 50.000. 

Volume  after  introducing  hydrogen 8o.b 

Volume  after  combination  of  oxygen  and  hydrogen  .  48.5 
Contraction  in  volume  due  to  combination (80  cc.  —48.5  cc.)  31 .5 
Volume  of  oxygen  in  50  cc.  of  air  (|  of  31.5)  .  .  .  .  10.5 


THE  ATMOSPHERE  87 

All  these  methods  agree  in  showing  that  100  volumes  of 
dry  air  contain  approximately  2 1  volumes  of  oxygen. 

2.  Determination  of  nitrogen.    If  the  gas  left  after  the 
removal  of  oxygen  from  a  portion  of  air  is  passed  over 
heated  magnesium,  the  nitrogen  is  with.dr.awn,  argon  and 
the  other  rare  elements  being  left.    It  may  thus  be  shown 
that  of  the  79  volumes  of  gas  left  after  the  removal  of  the 
oxygen  from  100  volumes  of  air,  approximately  78  are  nitro- 
gen and  0.93  argon.    The  other  elements  are  present  in 
such  small  quantities  that  they  may  be  neglected. 

3.  Determination  of  carbon  dioxide.    The  percentage  of 
carbon  dioxide  in  any  given  volume  of  air  may  be  deter- 
mined by  passing  the  air  over  calcium  hydroxide  or  some 
other  compound  which  will  combine  with  the  carbon  dioxide. 
The  increase  in  the  weight  of  the  hydroxide   equals  the 
weight  of  the  carbon  dioxide  absorbed.    The  amount  pres- 
ent in  the  open  normal  air  is  from  3  to  4  parts  by  volume 
in  10,000  volumes  of  air,  or  about  0.04%. 

4.  Determination  of  water  vapor.    The  water  vapor  pres- 
ent in  a  given  volume  of  air  may  be  determined  by  passing 
the  air  over  calcium   chloride  (or  some    other  compound 
which  has  a  strong  affinity  for  water),  and  noting  the  increase 
in  the  weight  of  the  chloride.    The  amount  present  varies 
not  only  with  the  locality,  but  there  is  a  wide  variation  from 
day  to  day  in  the  same  locality  because  of  the  winds  and 
changes  in  temperature. 

Processes  affecting  the  composition  of  the  air.  The  most 
important  of  these  processes  are  the  following. 

i .  Respiration.  In  the  process  of  respiration  some  of  the 
oxygen  in  the  inhaled  air  is  absorbed  by  the  blood  and  car- 
ried to  all  parts  of  the  body,  where  it  combines  with  the 
carbon  of  the  worn-out  tissues.  The  products  of  oxidation 


88       AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

are  carried  back  to  the  lungs  and  exhaled  in  the  form  of 
carbon  dioxide.  The  amount  exhaled  by  an  adult  averages 
about  20  1.  per  hour.  Hence  in  a  poorly  ventilated  room 
occupied  by  a  number  of  people  the  amount  of  carbon  diox- 
ide rapidly  increases.  While  this  gas  is  not  poisonous  unless 
present  in  large  amounts,  nevertheless  air  containing  more 
than  15  parts  in  10,000  is  not  fit  for  respiration. 

2.  Combustion.    All  of  the  ordinary  forms  of  fuel  con- 
tain large  percentages  of  carbon.    On  burning,  this  carbon 
combines  with  oxygen  in  the  air,  forming  carbon  dioxide. 
Combustion  and  respiration,  therefore,  tend  to  diminish  the 
amount  of  oxygen  in  the  air  and  to  increase  the  amount  of 
carbon  dioxide. 

3.  Action  of  plants.    Plants  have  the  power,  when  in  the 
sunlight,  of  absorbing  carbon  dioxide  from  the  air,  retaining 
the  carbon  and  returning  at  least  a  portion  of  the  oxygen 
to  the  air.     It  will  be  observed  that  these  changes  are  just 
the  opposite  of  those  brought  about  by  the  processes  of 
respiration  and  combustion. 

Poisonous  effect  of  exhaled  air.  The  differences  in  the  percent- 
ages of  oxygen,  carbon  dioxide,  and  moisture  present  in  inhaled  air 
and  exhaled  air  are  shown  in  the  following  analyses. 

INHALED  AIR      EXHALED  AIR 

Oxygen    . 21.00%  16.00% 

Carbon  dioxide 0.04  4.38 

Moisture variable         saturated 

The  foul  odor  of  respired  air  is  due  to  the  presence  of  a  certain 
amount  of  organic  matter.  It  is  possible  that  this  organic  matter 
rather  than  the  carbon  dioxide  is  responsible  for  the  injurious  effects 
which  follow  the  respiration  of  impure  air.  The  extent  of  such 
organic  impurities  present  may  be  judged,  however,  by  the  amount 
of  carbon  dioxide  present,  since  the  two  are  exhaled  together. 

The  cycle  of  carbon  in  nature.  Under  the  influence  of  sunlight,  the 
carbon  dioxide  absorbed  from  the  air  by  plants  reacts  with  water 


THE  ATMOSPHERE  89 

and  small  amounts  of  other  substances  absorbed  from  the  soil  to 
form  complex  compounds  of  carbon  which  constitute  the  essential 
part  of  the  plant  tissue.  This  reaction  is  attended  by  the  evolution 
of  oxygen,  which  is  restored  to  the  air.  The  compounds  resulting 
from  these  changes  are  much  richer  in  their  energy  content  than  are 
the  substances  from  which  they  are  formed  ;  hence  a  certain  amount 
of  energy  must  have  been  absorbed  in  their  formation.  The  source 
of  this  energy  is  the  sun's  rays. 

If  the  plant  is  burned,  the  changes  which  took  place  in  the  forma- 
tion of  the  compounds  present  are  largely  reversed.  The  carbon  and 
hydrogen  present  combine  with  oxygen  taken  from  the  air  to  form 
carbon  dioxide  and  water,  while  the  energy  absorbed  from  the  sun's 
rays  is  liberated  in  the  form  of  energy  of  heat.  If,  on  the  other  hand, 
the  plant  is  used  as  food,  the  compounds  present  are  used  in  building 
up  the  tissues  of  the  body.  When  this  tissue  breaks  down,  the 
changes  which  it  undergoes  are  very  similar  to  those  which  take 
place  when  the  plant  is  burned.  The  carbon  and  hydrogen  combine 
with  the  inhaled  oxygen  to  form  carbon  dioxide  and  water,  which 
are  exhaled.  The  energy  possessed  by  the  complex  substances  is 
liberated  partly  in  the  form  of  energy  of  heat,  which  maintains  the 
heat  of  the  body,  and  partly  in  the  various  forms  of  muscular  energy. 
The  carbon  originally  absorbed  from  the  air  by  the  plant  in  the  form 
of  carbon  dioxide  is  thus  restored  to  the  air  and  is  ready  to  repeat 
the  cycle  of  changes. 

The  composition  of  the  air  is  constant.  Notwithstanding 
the  changes  constantly  taking  place  which  tend  to  alter 
the  composition  of  the  air,  the  results  of  a  great  many 
analyses  of  air  collected  in  the  open  fields  show  that  the 
percentages  of  oxygen  and  nitrogen  as  well  as  of  carbon 
dioxide  are  very  nearly  constant.  Indeed,  so  constant  are 
the  percentages  of  oxygen  and  nitrogen  that  the  question 
has  arisen,  whether  these  two  elements  are  not  combined 
in  the-  air,  forming  a  definite  chemical  compound.  That 
the  two.  are  not  combined  but  are  simply  mixed  together 
can  be  shown  in  a  number  of  ways,  among  which  are  the 
following. 


90        AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

1.  When  air  dissolves  in  water  it  has  been  found  that 
the  ratio  of  oxygen  to  nitrogen  in  the  dissolved  air  is  no 
longer  21  :  78,  but  more  nearly  35  :  65.    If  it  were  a  chem- 
ical compound,  the  ratio  of  oxygen  to  nitrogen  would  not 
be  changed  by  solution  in  water. 

2.  A  chemical  compound  in  the  form  of  a  liquid  has  a 
definite  boiling  point.    Water,  for  example,  boils  at   100°. 
Moreover  the  steam  which  is  thus  formed  has  the  same 
composition  as  the  water.    The  boiling  point  of  liquid  air, 
on  the  other  hand,  gradually  rises  as  the  liquid  boils,  the 
nitrogen  escaping  first  followed  by  the  oxygen.    If  the  two 
were  combined,  they  would  pass  off  together  in  the  ratio  in 
which  they  are  found  in  the  air. 

Why  the  air  has  a  constant  composition.  If  air  is  a  mix- 
ture and  changes  are  constantly  taking  place  which  tend 
to  modify  its  composition,  how,  then,  do  we  account  for  the 
constancy  of  composition  which  the  analyses  reveal  ?  This 
is  explained  by  several  facts,  (i)  The  changes  which  are 
caused  by  the  processes  of  combustion  and  respiration,  on 
the  one  hand,  and  the  action  of  plants,  on  the  other,  tend  to 
equalize  each  other.  (2)  The  winds  keep  the  air  in  con- 
stant motion  and  so  prevent  local  changes.  (3)  The  vol- 
ume of  the  air  is  so  vast  and  the  changes  which  occur  are  so 
small  compared  with  the  total  amount  of  air  that  they  cannot 
be  readily  detected.  (4)  Finally  it  must  be  noted  that  only 
air  collected  in  the  open  fields  shows  this  constancy  in 
composition.  The  air  in  a  poorly  ventilated  room  occupied 
by  a  number  of  people  rapidly  changes  in  composition. 

The  properties  of  the  air.  Inasmuch  as  air  is  com- 
posed principally  of  a  mixture  of  oxygen  and  nitrogen, 
which  elements  have  already  been  discussed,  its  proper- 
ties may  be  inferred  largely  from  those  of  the  two  gases. 


THE  ATMOSPHERE  91 

One  liter  weighs  1.2923  g.  It  is  thus  14.38  times  as  heavy 
as  hydrogen.  At  the  sea  level  it  exerts  an  average  pressure 
sufficient  to  sustain  a  column  of  mercury  760  mm.  in  height. 
This  is  taken  as  the  standard  pressure  in  determining  the 
volumes  of  gases  as  well  as  the  boiling  points  of  liquids. 
Water  may  be  made  to  boil  at  any  temperature  between  o° 
and  considerably  above  100°  by  simply  varying  the  pressure. 
It  is  only  when  the  pressure  upon  it  is  equal  to  the  normal 
pressure  of  the  atmosphere  at  the  sea  level,  as  indicated  by 
a  barometric  reading  of  760  mm.,  that  it  boils  at  100°. 

Preparation  of  liquid  air.  Attention  has  been  called  to 
the  fact  that  both  oxygen  and  nitrogen  can  be  obtained  in 
the  liquid  state  by  strongly  cooling  the  gases  and  applying 
great  pressure  to  them.  Since  air  is  largely  a  mixture  of 
these  two  gases,  it  can  be  liquefied  by  the  same  methods. 

The  methods  for  liquefying  air  have  been  simplified  greatly  in 
that  the  low  temperature  required. is  obtained  by  allowing  a  portion 
of  the  compressed  air  to  expand.  The  expansion 
of  a  gas  is  always  attended  by  the  absorption  of 
heat.  In  liquefying  air  the  apparatus  is  so  con- 
structed that  the  heat  absorbed  is  withdrawn 
from  air  already  under  great  pressure.^  This 
process  is  continued  until  the  temperature  is 
lowered  to  the  point  oTliquefaction. 

The  Dewar  bulb.  It  is  not  possible  to, 
preserve  air  in  the  liquid  state  in  a  closed 
vessel,  on  account  of  the  enormous  pres- 
sure exerted  by  it  in  its  tendency  to  pass 
into  the  gaseous  state.  It  may  however  be  preserved 
for  some  hours  or  even  days  before  it  will  completely 
evaporate,  by  simply  placing  it  in  an  open  vessel  sur- 
rounded by  a  nonconducting  material.  The  most  efficient 
vessel  for  this  purpose  is  the  Dewar  bulb  shown  in  Fig.  29. 


92        AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

The  air  is  withdrawn  from  the  space  between  the  two 
walls,  thus  making  it  nonconducting. 

Properties  and  uses  of  liquid  air.  When  first  prepared, 
liquid  air  is  cloudy  because  of  the  presence  of  particles  of 
solid  carbon  dioxide.  These  may  be  filtered  off,  leaving 
a  liquid  of  slightly  bluish  color.  It  begins  to  boil  at  about 
—  190°,  the  nitrogen  passing  off  first,  gradually  followed 
by  the  oxygen,  the  last  portions  being  nearly  pure  oxygen. 
To  a  certain  extent  oxygen  is  now  prepared  in  this  way  for 
commercial  purposes. 

The  extremely  low  temperature  of  liquid  air  may  be 
inferred  from  the  fact  that  mercury  when  cooled  by  it  is 
frozen  to  a  mass  so  hard  that  it  may  be  used  for  driving 
nails. 

Liquid  air  is  used  in  the  preparation  of  oxygen  and  as 
a  cooling  agent  in  the  study  of  the  properties  of  matter  at 
low  temperatures.  It  has  thus  been  found  that  elements 
at  extremely  low  temperatures  largely  lose  their  chemical 
activity. 

EXERCISES 

1.  When  oxygen  and  nitrogen  are  mixed  in  the  proportion  in 
which    they  exist  in  the  atmosphere,   heat   is  neither  evolved  nor 
absorbed  by  the  process.    What  important  point  does  this  suggest  ? 

2.  What  essential  constituent  of  the  air  is  found  in  larger  amount 
in  manufacturing  districts  than  in  the  open  country  ? 

3.  Can  you  suggest  any  reason  why  the  growth  of  clover  in  a  field 
improves  the  soil  ? 

4.  Why  are  the  inner  walls  of  a  Dewar  bulb  sometimes  coated 
with  a  film  of  silver  ? 

5.  To  what  is  the  blue  color  of  liquid  air  due?    Does  this  color 
increase  in  intensity  on  standing  ? 

6.  When  ice  is  placed  in  a  vessel  containing  liquid  air,  the  latter 
boils  violently.    Explain. 


THE  ATMOSPHERE  93 

7.  Taking  the  volumes  of  the  oxygen  and  nitrogen  in  100  volumes 
of  air  as  21  and  78  respectively,  calculate  the  percentages  of  these 
elements  present  by  weight. 

8.  Would  combustion  be  more  intense  in  liquid  air  than  in  the 
gaseous  substance? 

9.  A   tube   containing  calcium   chloride   was    found  to   weigh 
30. 1293  g.    A  volume  of  air  which  weighed   15. 2134  g.  was  passed 
through,  after  which  the  weight  of  the  tube  was  found  to  be  30.3405  g. 
What  was  the  percentage  amount  of  moisture  present  in  the  air  ? 

10.  10  1.  of  air  measured  at  20°  and  740  mm.  passed  through 
lime  water  caused  the  precipitation  of  0.0102  g.  of  CaCO3.  Find  the 
number  of  volumes  of  carbon  dioxide  in  10,000  volumes  of  the  air. 


CHAPTER   IX 


SOLUTIONS 

Definitions.  When  a  substance  disappears  in  a  liquid  in 
such  a  way  as  to  thoroughly  mix  with  it  and  to  be  lost  to 
sight  as  an  individual  body,  the  resulting  liquid  is  called  a 
solution.  The  liquid  in  which  the  substance  dissolves  is 
called  the  solvent,  while  the  dissolved  substance  is  called 
the  solute. 

Classes  of  solutions.  Matter  in  any  one  of  its  physical 
states  may  dissolve  in  a  liquid,  so  that  we  may  have  solu- 
tions of  gases,  of  liquids,  and  of  solids.  Solutions  of  liquids 
in  liquids  are  not  often  mentioned  in  the  following  pages, 
but  the  other  two  classes  will  become  very  familiar  in  the 
course  of  our  study,  and  deserve  special  attention. 

SOLUTION  OF  GASES  IN  LIQUIDS 

It  has  already  been  stated  that  oxygen,  hydrogen,  and 
nitrogen  are  slightly  soluble  in  water.  Accurate  study  has 

led  to  the  concision  that 
all  gases  arex^oluble  to 
some  Qnlent  not  only  in 
water  but  in  many  other 
liquids.  The  amount  of  a 
gas  which  will  dissolve  in  a 
liquid  depends  upon  a  num- 
ber of  conditions,  and  these 
FlG<30  can  best  be  understood  by 

94 


SOLUTIONS 


95 


supposing  a  vessel  B  (Fig.  30),  to  be  filled  with  the  gas 
and  inverted  over  the  liquid.  Under  these  circumstances 
the  gas  cannot  escape  or  become  mixed  with  another  gas. 

Circumstances  affecting  the  solubility  of  gases.  A  num- 
ber of  circumstances  affect  the  solubility  of  a  gas  in  a  liquid. 

1 .  Nature  of  the  gas.    Other  conditions  being  equal,  each 
gas  has  its  own  peculiar  solubility,  just  as  it  has  its  own 
special  taste  or  odor.    The  solubility  of  gases  varies  between 
wide  limits,  as  will  be  seen  from  the  following  table,  but  as 
a  rule  a  given  volume  of  a  liquid  will  not  dissolve  more  than 
two  or  three  times  its  own  volume  of  a  gas. 

Solubility  of  Gases  in  Water 
i  1.  of  water  at  760  mm.  pressure  and  at  o°  will  dissolve  : 

Ammonia 1148.00  1. 

Hydrochloric  acid 503.00 

Sulphur  dioxide .  79-79 

Carbon  dioxide i  .80 

Oxygen 41.14  cc. 

Hydrogen 21.15 

Nitrogen        20.03 

In  the  case  of  very  soluble  gases,  such  as  the  first  three 
in  the  table,  it  is  probable  that  chemical  combination 
between  the  liquid  and  the  gas  takes  place. 

2 .  Nature  of  the  liquid.    The  character"  of  the  liquid  has 
much  influence  upon  the  solubility  of  a  gas.    Water,  alco- 
hol, and  ether  have  each  its  own  peculiar  solvent  power. 
From  the  solubility  of  a  gas  in  water,  no  prediction  can  be 
made  as  to  its  solubility  in  other  liquids. 

3.  Influence  of  pressure.    It   has  been  found  that  the 
weight  of  gas  which  dissolves  in  a  given  case  is  propor- 
tional   to    the    pressure    exerted    upon    the    gas.     If    the 


96        AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

pressure  is  doubled,  the  weight  of  gas  going  into  solution 
is  doubled ;  if  the  pressure  is  diminished  to  one  half  of  its 
original  value,  half  of  the  dissolved  gas  will  escape.  Under 
high  pressure,  large  quantities  of  gas  can  be  dissolved  in  a 
liquid,  and  when  the  pressure  is  removed  the  gas  escapes, 
causing  the  liquid  to  foam  or  effervesce. 

4.  Influence  of  temperature.  In  general,  the  lower  the 
temperature  of  the  liquid,  the  larger  the  quantity  of  gas 
which  it  can  dissolve.  1000  volumes  of  water  at  o°  will 
dissolve  41.14  volumes  of  oxygen  ;  at  50°,  18.37  volumes  ; 
at  1 00°  none  at  all.  While  most  gases  can  be  expelled  from 
a  liquid  by  boiling  the  solution,  some  cannot.  For  example, 
it  is  not  possible  to  expel  hydrochloric  acid  gas  completely 
from  its  solution  by  boiling. 

SOLUTION  OF  SOLIDS  IN  LIQUIDS 

This  is  the  most  familiar  class  of  solutions,  since  in  the 
laboratory  substances  are  much  more  frequently  used  in  the 
form  of  solutions  than  in  the  solid  state. 

Circumstances  affecting  the  solubility  of  a  solid.  The 
solubility  of  a  solid  in  a  liquid  depends  upon  several  factors. 

I .  Nature  of  the  solid.  Other  conditions  being  the  same, 
solids  vary  greatly  in  their  solubility  in  liquids.  This  is 
illustrated  in  the  following  table  : 

Table  of  Solubility  of  Solids  at  18° 
100  cc.  of  water  will  dissolve : 

Calcium  chloride 7i.og. 

Sodium  chloride 35.9 

Potassium  nitrate 29.1 

Copper  sulphate 21.4 

Calcium  sulphate 0.207- 


SOLUTIONS  97 

No  solids  are  absolutely  insoluble,  but  the  amount  dis- 
solved may  be  so  small  as  to  be  of  no  significance  for  most 
purposes.  Thus  barium  sulphate,  one  of  the  most  insoluble 
of  common  substances,  dissolves  in  water  to  the  extent  of 
i  part  in  400,000. 

2.  Nature  of  the  solvent.     Liquids  vary  much  in  their 
power  to  dissolve  solids.    Some  are  said  to  be  good  solvents, 
since  they  dissolve  a  great  variety  of  substances  and  consid- 
erable quantities  of  them.    Others  have  small  solvent  power, 
dissolving  few  substances,  and  those  to  a  slight  extent  only. 
Broadly  speaking,  water  is  the  most  general  solvent,  and 
alcohol  is  perhaps  second  in  solvent  power. 

3.  Temperature.    The  weight  of  a  solid  which  a  given 
liquid  can  dissolve  varies  with  the  temperature.    Usually  it 
increases  rapidly  as  the  temperature  rises,  so  that  the  boil- 
ing liquid  dissolves  several  times  the  weight  which  the  cold 
liquid  will  dissolve.    In  some  instances,  as  in  the  case  of 
common  salt  dissolved  in  water,  the  temperature  has  little 
influence  upon  the  solubility,  and  a  few  solids  are  more 
soluble  in  cold  water  than  in  hot.    The  following  examples 
will  serve  as  illustrations  : 

Table  of  Solubility  at  O°  and  at  IOO° 

100  cc.  of  water  will  dissolve : 

f 

At  o°   _  At  100° 

Calcium  chloride 49.6  g.  155.0  g. 

Sodium  chloride 35.7  39.8  ,. 

Potassium  nitrate 13.3  247.-° 

Copper  sulphate 15.5  73-5. 

Calcium  sulphate 0.20  c  0.217 

Calcium  hydroxide -0.173  0.079 

Saturated  solutions.    A  liquid  will  not  dissolve  an  unlimited  quan- 
tity of  a  solid.    On  adding  the  solid  to  the  liquid  in  small  portions  at 


98        AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

a  time,  it  will  be  found  that  a  point  is  reached  at  which  the  liquid  will 
not  dissolve  more  of  the  solid  at  that  temperature.  The  solid  and  the 
solution  remain  in  contact  with  each  other  unchanged.  This  condi- 
tion may  be  described  by  saying  that  they  are  in  equilibrium  with 
each  other.  A  solution  is  said  to  be  saturated  when  it  remains 
unchanged  in  concentration  in  contact  with  some  of  the  solid.  The 
weight  of  the  solid  which  will  completely  saturate  a  definite  volume 
of  a  liquid  at  a  given  temperature  is  called  the  solubility  of  the  sub- 
stance at  that  temperature. 

Supersaturated  solutions.  When  a  solution,  saturated  at  a  given 
temperature,  is  allowed  to  cool  it  sometimes  happens  that  no  solid  crys- 
tallizes out.  This  is  very  likely  to  occur  when  the  vessel  used  is 
perfectly  smooth  and  the  solution  is  not  disturbed  in  any  way.  Such 
a  solution  is  said  to  be  supersaturated.  That  this  condition  is  unstable 
can  be  shown  by  adding  a  crystal  of  the  solid  to  the  solution.  All  of 
the  solid  in  excess  of  the  quantity  required  to  saturate  the  solution  at 
this  temperature  will  at  once  crystallize  out,  leaving  the  solution  satu- 
rated. Supersaturation  may  also  be  overcome  in  many  cases  by  vigor- 
ously shaking  or  stirring  the  solution. 

General  physical  properties  of  solutions.  A  few  general 
statements  may  be  made  in  reference  to  the  physical  prop- 
erties of  solutions. 

1 .  Distribution  of  the  solid  in  the  liquid.    A  solid,  when 
dissolved,  tends  to  distribute  itself  uniformly  through  the 
liquid,   so  that  every  part  of  the   solution  has   the   same 
concentration.    The  process  goes    on   very  slowly   unless 
hastened  by  stirring  or  shaking  the  solution.    Thus,   if  a 
few  crystals  of  a  highly  colored  substance  such  as  copper  sul- 
phate are  placed  in  the  bottom  of  a  tall  vessel  full  of  water, 
it  will  take  weeks  for  the  solution  to  become  uniformly 
colored. 

2.  Boiling  points  of  solutions.    The  boiling  point  of  a 
liquid  is  raised  by  the  presence  of  a  substance  dissolved  in 
it.    In  general  the  extent  to  which  the  boiling  point  of  a 
solvent  is  raised  by  a  given  substance  is  proportional  to  the 


SOLUTIONS 


99 


concentration  of  the  solution,  that  is,  to  the  weight  of  the 
substance  dissolved  in  a  definite  weight  of  the  solvent. 

3.  Freezing  points  of  solutions.  A  solution  freezes  at  a 
lower  temperature  than  ,the  pure  solvent.  The  lowering  of 
the  freezing  point  obeys  the  same  law  which  holds  for  the 
raising  of  the  boiling  point  :  the  extent  of  lowering  is  pro- 
portional to  the  weight  of  dissolved  substance,  that  is,  to  the 
concentration  of  the  solution. 

Electrolysis  of  solutions.  Pure  water  does  not  appreciably 
conduct  the  electric  current.  If,  however,  certain  substances 
such  as  common  salt  are  dissolved  in  the  water,  the  result- 
ing solutions  are  found  to  be  conductors  of  electricity. 
Such  solutions  are  called  electro- 
lytes. When  the  current  passes 
through  an  electrolyte  some  chemi- 
cal change  always  takes  place.  This 
change  is  called  electrolysis. 
^  The  general  method  used  in  the 
electrolysis  of  a  solution  is  illus- 
trated in  Fig.  31.  The  vessel  D  contains  the  electrolyte. 
Two  plates  or  rods,  A  and  B,  made  of  suitable  material, 
are  connected  with  the  wires  from  a  battery  (or  dynamo) 
and  dipped  into  the  electrolyte,  as  shown  in  the  figure. 
These  plates  or  rods  are  called  electrodes.  The  electrode 
connected  with  the  zinc  plate  of  the  battery  is  the  negative 
electrode  or  cathode,  while  that  connected  with  the  carbon 
plate  is  the  positive  electrode  or  anode.  "+~" 

Theory  of  electrolytic  dissociation.  The  facts  which  have 
just  been  described  in  connection  with  solutions,  together 
with  many  others,  have  led  chemists  to  adopt  a  theory  of 
solutions  called  the  theory  of  electrolytic  dissociation^  .  The 
main  assumptions  in  this  theory  are  the  following. 


loo     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

i  .  Formation  of  ions.  Many  compounds  when  dissolved 
in  water  undergo  an  important  change.  A  portion  of  their 
molecules  fall  apart,  or  dissociate,  into  two  or  more  parts, 
called  ions.  Thus  sodium  nitrate  (NaNO3)  dissociates  into 
the  ions  Na  and  NO3  ;  sodium  chloride,  into  the  ions  Na 
and  Cl.  These  ions  are  free  to  move  about  in  the  solution 
independently  of  each  other  like  independent  molecules,  and 
for  this  reason  were  given  the  name  ion,  which  signifies  a 
wanderer. 

2.  The   electrical  charge   of  ions.    Each    ion   carries    a 
heavy  electrical  charge,  and  in  this  respect  differs  from  an 
atom  or  molecule.    It  is  evident  that  the  sodium  in  the  form 
of  an  ion  must  differ  in  some  important  way  from  ordinary 
sodium,  for  sodium  ions,  formed  from  sodium  nitrate,  give 
no  visible  evidence  of  their  presence  in   water,   whereas 
metallic  sodium  at  once  decomposes  the  water.    The  elec- 
trical charge,  therefore,  greatly  modifies  the  usual  chemical 
properties  of  the  element. 

3.  The  positive  charges  equal  the  negative  charges.    The 
ions  formed  by  the  dissociation  of  any  molecule  are  of  two 
kinds.    One  kind  is  charged  with  positive  electricity  and 
the  other  with  negative  electricity  ;  moreover  the  sum  of  all 
the  positive  charges  is  always  equal  to  the  sum  of  all  the 
negative  charges.    The  solution  as  a  whole  is  therefore  elec- 
trically neutral.    If  we  represent  dissociation  by  the  usual 
chemical  equations,  with  the  electrical  charges  indicated  by 
+  and  —  signs  following  the  symbols,  the  dissociation  of 
sodium  chloride  molecules  is  represented  thus  : 


The  positive  charge  on  each  sodium  ion  exactly  equals  the 
negative  charge  on  each  chlorine  ion. 


SOLUTIONS  101 

Sodium  sulphate  dissociates,  as  shown  in  the  equation 
^2Na+,  SO4~ 


Here  the  positive  charge  on  the  two  sodium  ions  equals 
the  double  negative  charge  on  the  SO4  ion. 

4.  Not  all  compounds  dissociate.    Only  those  compounds  ] 
dissociate  whose  solutions  form  electrolytes.    Thus  salt  dis-  1 
sociates  when  dissolved  in  water,   the    resulting   solution    f 
being  an  electrolyte.    Sugar,  on  the  other  hand,  does  not   / 
dissociate  and  its  solution  is  not  a  conductor  of  the  electric/ 
current. 

5.  Extent   of  dissociation   differs   in    different  liquids. 
While    compounds    most    readily    undergo   dissociation    in 
water,  yet  dissociation  often  occurs  to  a  limited  extent  when 
solution  takes  place  in  liquids  other  than  water.    In  the 
discussion  of  solutions  it  will  be  understood  that  the  solvent 
is  water  unless  otherwise  noted. 

The  theory  of  electrolytic  dissociation  and  the  properties 
of  solutions.  In  order  to  be  of  value,  this  theory  must  give 
a  reasonable  explanation  of  the  properties  of  solutions.  Let 
us  now  see  if  the  theory  is  in  harmony  with  certain  of  these 
properties. 

The  theory  of  electrolytic  dissociation  and  the  boiling 
and  freezing  points  of  solutions.  We  have  seen  that  the 
boiling  point  of  a  solution  of  a  substance  is  raised  in  pro- 
portion to  the  concentration  of  the  dissolved  substance. 
This  is  but  another  way  of  saying  that  the  change  in  the 
boiling  point  of  the  solution  is  proportional  to  the  number 
of  molecules  of  the  dissolved  substance  present  in  the 
solution. 

It  has  been  found,  however,  that  in  the  case  of  electro- 
lytes the  boiling  point  is  raised  more  than  it  should  be  to 

;  ::\:5Vr 
»«••»»  •        ^ 


102     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

conform  to  this  law.  If  the  solute  dissociates  into  ions,  the 
reason  for  this  becomes  clear.  Each  ion  has  the  same  effect 
on  the  boiling  point  as  a  molecule,  and  since  their  number  is 
greater  than  the  number  of  molecules  from  which  they  were 
formed,  the  effect  on  the  boiling  point  is  abnormally  great. 

In  a  similar  way,  the  theory  furnishes  an  explanation  of 
the  abnormal  lowering  of  the  freezing  point  of  electrolytes. 

The  theory  of  electrolytic  dissociation  and  electrolysis. 
The  changes  taking  place  during  electrolysis  harmonize  very 
completely  with  the  theory  of  dissociation.  This  will  become 
clear  from  a  study  of  the  following  examples. 

i.  Electrolysis  of  sodium  chloride.  Fig.  32  represents  a 
vessel  in  which  the  electrolyte  is  a  solution  of  sodium  chlo- 

ride  (NaCl). 
According  to 

,     the  dissociation 
~*-Na  Na  Na  Na  Na  d 

theory  the 
ci   ci  a   a   ci-+  -    '     n  , 

molecules  of 


sodium    chlo- 
ride dissociate 

into  the  ions  Na+  and  Cl~.  The  Na+  ions  are  attracted  to 
the  cathode  owing  to  its  large  negative  charge.  On  com- 
ing into  contact  with  the  cathode,  the  Na+  ions  give  up 
their  positive  charge  and  are  then  ordinary  sodium  atoms.1 
They  immediately  decompose  the  water  according  to  the 

equation 

Na  +  H20  =  NaOH  +  H, 

and  hydrogen  is  evolved  about  the  cathode. 

The  chlorine  ions  on  being  discharged  at  the  anode  in 
similar  manner  may  either  be  given  off  as  chlorine  gas,  or 
may  attack  the  water,  as  represented  in  the  equation 

O. 


SOLUTIONS 


103 


•  2 .  Electrolysis  of  water.  The  reason  for  the  addition  of 
sulphuric  acid  to  water  in  the  preparation  of  oxygen  and 
hydrogen  by  electrolysis  will  now  be  clear.  Water  itself  is 
not  an  electrolyte  to  an  appreciable  extent ;  that  is,  it  does 
not  form  enough  ions  to  carry  a  current.  Sulphuric  acid  dis- 
solved in  water  is  an  electrolyte,  and  dissociates  into  the  ions 
2H+  and  SO^".  In  the  process  of  electrolysis  of  the  solu- 
tion, the  hydrogen  ions  travel  to  the  cathode,  and  on  being 
discharged  escape  as  hydrogen  gas.  The  SO4  ions,  wheft 
discharged  at  the  anode,  act  upon  water,  setting  free  oxy- 
gen and  once  more  forming  sulphuric  acid  : 

S04  +  H20  =  H2S04  +  O. 

The  sulphuric  acid  can  again  dissociate  and  the  process 
repeat  itself  as  long  as  any  water  is  left.  Hence  the  hydro- 
gen and  oxygen  set  free  in  the  electrolysis  of  water  really 
come  directly  from  the  acid  but  indirectly  from  the  water. 
3.  Electrolysis  of  sodium  sulphate.  In  a  similar  way, 
sodium  sulphate  (Na2SO4),  when  in  solution,  gives  the  ions 
2Na+  and  SO4~.  On  being  discharged,  the  sodium  atoms 
decompose  water  about  the  cathode,  _  + 

as  in  the  case  of  sodium  chloride, 
while  the  SO4  ions  when  discharged 
at  the  anode  decompose  the  water,  as 
represented  in  the  equation 

SO4  +  H2O  =  H2SO4  +  O 

That  new  substances  are  formed  at 
the  cathode  and  anode  may  be  shown 
in  the  following  way.  A  U-tube,  such 

FIG. 33 

as  is  represented  in  Fig.  33,  is  par- 
tially filled  with  a  solution  of  sodium  sulphate,  and  the  liquid 
in  one  arm  is  colored  with  red  litmus,  that  in  the  other 


104     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

with  blue  litmus.  An  electrode  placed  in  the  red  solution 
is  made  to  serve  as  cathode,  while  one  in  the  blue  solution 
is  made  the  anode.  On  allowing  the  current  to  pass,  the 
blue  solution  turns  red,  while  the  red  solution  turns  blue. 
These  are  exactly  the  changes  which  would  take  place  if 
sodium  hydroxide  and  sulphuric  acid  were  to  be  set  free  at 
the  electrodes,  as  required  by  the  theory. 

The  properties  of  electrolytes  depend  upon  the  ions  pres- 
ent. When  a  substance  capable  of  dissociating  into  ions  is 
dissolved  in  water,  the  properties  of  the  solution  will  depend 
upon  two  factors  :  (i)  the  ions  formed  from  the  substance; 
(2)  the  undissociated  molecules.  Since  the  ions  are  usu- 
ally more  active  chemically  than  the  molecules,  most  of  the 
chemical  properties  of  an  electrolyte  are  due  to  the  ions 
rather  than  to  the  molecules. 

The  solutions  of  any  two  substances  which  give  the  same 
ion  will  have  certain  properties  in  common.  Thus  all  solu- 
tions containing  the  copper  ion  (Cu++)  are  blue,  unless  the 
color  is  modified  by  the  presence  of  ions  or  molecules  hav- 
ing some  other  color. 

EXERCISES 

1.  Distinguish  clearly  between  the  following  terms :  electrolysis, 
electrolyte,   electrolytic  dissociation,   ions,   solute,   solvent,  solution, 
saturated  solution,  and  supersaturated  solution. 

2.  Why  does  the  water  from  some  natural  springs  effervesce  ? 

3.  (cT)  Why  does  not  the  water  of  the  ocean  freeze  ?    (^)  Why 
will  ice  and  salt  produce  a  lower  temperature  than  ice  alone  ? 

4.  Why  does  shaking   or  stirring  make   a  solid  dissolve  more 
rapidly  in  a  liquid? 

5.  By  experiment  it  was  found  that  a  certain  volume  of  water 
was  saturated  at  100°  with  114  g.  of  potassium  nitrate.    On  cooling 
to  o°  a  portion  of  the  substance  crystallized.    (#)  How  many  grams 
of  the  substance  remained  in  solution?    (£)  What  was  the  strength 


SOLUTIONS  105 

of  the    solution    at   18°?    (V)  How  much   water  had  been  used  in 
the  experiment  ? 

6.  (#)  10  g.  of  common  salt  were  dissolved   in  water  and  the 
solution  evaporated   to    dryness  ;    what  weight  of   solid  was  left  ? 
(£)  10  g.  of  zinc  were  dissolved  in  hydrochloric  acid  and  the  solu- 
tion evaporated  to  dryness  ;  what  weight  of  solid  was  left  ? 

7.  Account   for   the   fact   that  sugar  sometimes  deposits   from 
molasses,  even  when  no  evaporation  has  taken  place. 

8.  (fl)  From  the  standpoint  of  the  theory  of  electrolytic  dissocia- 
tion, write  the  simple  equation  for  a  dilute  solution  of  copper  sulphate 
(CuSO4);  this  solution  is  blue.    (£)  In  the  same  manner,  write  one 
for  sodium  sulphate  ;  this  solution  is  colorless,    (c)  How  would  you 
account  for  the  color  of  the  copper  sulphate  solution  ? 

9.  (a)  As  in  the  preceding  exercise,  write  a  simple  equation  for 
a  dilute  solution  of  copper  chloride  (CuCl2);  this  solution  is  blue. 
(£)  In  the  same  manner,  write  one  for  sodium  chloride ;  this  solution 
is  colorless.    To  what  is  the  blue  color  due  ? 

10.  What  component  is  present  in  concentrated  sulphuric  acid 
that  is  almost  wanting  in  very  dilute  sulphuric  acid? 

11.  Why  will  vegetables  cook  faster  when  boiled  in  strong  salt 
water  than  when  boiled  in  pure  water? 

12.  How  do  you  explain  the  foaming  of  soda  water? 


CHAPTER  X 
ACIDS,  BASES,  AND  SALTS;   NEUTRALIZATION 

Acids,  bases,  and  salts.  The  three  classes  of  compounds 
known  respectively  as  acids,  bases,  and  salts  include  the 
great  majority  of  the  compounds  with  which  we  shall  have 
to  deal.  It  is  important,  therefore,  for  us  to  consider  each 
of  these  classes  in  a  systematic  way.  The  individual  mem- 
bers belonging  to  each  class  will  be  discussed  in  detail 
in  the  appropriate  places,  but  a  few  representatives  of  each 
class  will  be  described  in  this  chapter  with  special  refer- 
ence to  the  common  properties  in  accordance  with  which 
they  are  classified. 

The  familiar  acids.  Hydrochloric  acid  is  a  gas  composed 
of  hydrogen  and  chlorine,  and  has  the  formula  HC1.  The 
substance  is  very  soluble  in  water,  and  it  is  this  solution 
which  is  usually  called  hydrochloric  acid.  Nittic  acid  is  a 
liquid  composed  of  hydrogen,  nitrogen,  and  oxygen,-  having 
the  formula  HNO3.  As  sold  commercially  it  is  mixed  with 
about  32%  of  water.  SulpJmric  acid,  whose  composition  is 
represented  by  the  formula  H2SO4,  is  an  oily  liquid  nearjy 
twice  as  heavy  as  water,  and  is  commonly  called  oil  of  vitriol. 

Characteristics  of  acids,  (i)  All  acids  contain  hydrogen.  „ 
(2)  When  dissolved  in  water  the  molecules  of  the  acid  dis- 
sociate into  two  kinds  of  ions.  One  of  these  is  always  hydro- 
gen and  is  the  cation  (+),  while  the  other  consists  of  the 
remainder  of  the  molecule  and  is  the  anion  (— ).  (3)  The 
solution  tastes  sour.  (4)  It  has  the  power  to  change  the 

1 06 


ACIDS,  BASES,  AND  SALTS         107 

color  of  certain  substances  called  indicators.  Thus  blue 
litmus  is  changed  to  red,  and  yellow  methyl  orange  is 
changed  to  red.  Since  all  acids  produce  hydrogen  canons, 
while  the  anions  of  each  are-  different,  the  properties 
which  all  acids  have  in  common  when  in  solution,  such 
as  taste  and  action  on  indicators,  must  be  attributed  to 
the  hydrogen  ions.  __ 

DEFINITION  :  An  acid  is  a  substance  which  produces 
hydrogen  ions  when  dissolved  in  water  or  other  dissociat- 
ing liquids. 

Undissociated  acids.  When  acids  are  perfectly  free  from 
water,  or  are  dissolved  in  liquids  like  benzene  which  do 
not  have  the  power  of  'dissociating  them  into  ions,  they 
should  have  no  real  acid  properties.  This  is  found  to  be 
the  case.  Under  these  circumstaTfces  they  do  not  affect 
the  color  of  indicators  or  have  any  of  the  properties 
characteristic  of  acids. 

The  familiar  bases.  The  bases  most  used  in  the  labora- 
tory are  sodium  hydroxide  (NaOH),  potassium  hydroxide 
(KOH),  and  calcium  hydroxide  (Ca(OH)2).  These  are  white 
solids,  soluble  in  water,  the  latter  sparingly  so.  Some  bases 
are  vWy  difficultly  soluble  in  water.  The  very  soluble  ones 
with  most  pronounced  basic  properties  are  sometimes  called 
the  alkalis. 

Characteristics  of  bases,  (i)  All  bases  contain  hydrogen 
and  oxygen.  (2)  When  dissolved  in  water  the  molecules  of 
the  base  dissociate  into  two  kinds  of  ions.  One  of  these  is 
always  composed  of  oxygen  and  hydrogen  and  is  the  anion. 
It  has  the  formula  OH  and  is  called  the  hydroxyl  ion.  The 
remainder  of  the  molecule,  which  usually  consists  of  a 
single  atom,  is  the  cation,  (3)  The  solution  of  a  base  has 


108     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

a  soapy  feel  and  a  brackish  taste.  (4)  It  reverses  the  color 
change  produced  in  indicators  by  acids,  turning  red  litmus 
blue,  and  red  methyl  orange  yellow.  Since  all  bases  produce 
hydroxyl  anions,  while  the  cations  of  each  are  different,  the 
properties  which  all  bases  have  in  common  when  in  solution 
must  be  due  to  the  hydroxyl  ions. 

DEFINITION  :  A  base  is  a  substance  which  produces 
hydroxyl  ions  when  dissolved  in  water  or  other  dissociat- 
ing liquids. 

Undissociated  bases.  Bases,  in  the  absence  of  water  or 
when  dissolved  in  liquids  which  do  not  dissociate  them, 
should  have  none  of  the  properties  characteristic  of  this 
class  of  substances.  This  has  been  found  to  be  the  case. 
For  example,  they  have  no  effect  upon  indicators  under 
these  circumstances. 

Neutralization.  When  an  acid  and  a  base  are  brought 
together  in  solution  in  proper  proportion,  the  characteristic 
properties  of  each  disappear.  The  solution  tastes  neither 
sour  nor  brackish  ;  it  has  no  effect  upon  indicators.  There 
can  therefore  be  neither  hydrogen  nor  hydroxyl  ions  present 
in  the  solution.  A  study  of  reactions  of  this  kind  has  shown 
that  the  hydrogen  ions  of  the  acid  combine  with  the  hydroxyl 
ions  of  the  base  to  form  molecules  of  water,  water  being  a 
substance  which  is  not  appreciable?  dissociated  into  ions. 
This  action  of  an  acid  on  a  base  is  called  neutralization. 
The  following  equations  express  the  neutralization  of  the 
three  acids  by  three  bases,  water  being  formed  in  each  case. 

Na+,  OH-  +  H+,^C1-  =  Na+,  Cl~  +  H2O. 
K+,  OH-  +  H+,  NO^  =  K+,  NO^  +  H2O. 
Ca++,  (OH)-- +  H2++,  SO^-=  Ca++,  SO;~  +  2  H2O. 


ACIDS,  BASES,  AND  SALTS  1 09 

DEFINITION  :  (Neutralization  consists  in  the  union  of  the 
hydrogen  ion  of  an  acid  with  the  hydroxyl  ion  of  a  base  to 
form  water. 

Salts.  It  will  be  noticed  that  in  neutralization  the  anion 
of  the  acid  and  the  cation  of  the  base  are  not  changed.  If, 
however,  the  water  is  expelled  by  evaporation,  these  two 
ions  slowly  unite,  and  when  the  water  becomes  saturated 
with  the  substance  so  produced,  it  separates  in  the  form 
of  a  solid  called  a  salt. 

DEFINITION  :  A  salt  is  a  substance  formed  by  the  union 
of  the  amon  of  an  acid  with  the  cation  of  a  base. 

Characteristics  of  salts,  (i)  From  the  definition  of  a 
salt  it  will  be  seen  that  there  is  no  element  or  group  of 
elements  which  characterize  salts.  (2)  Salts  as  a  class  have 
no  peculiar  taste.  (3)  In  the  absence  of  all  other  substances 
they  are  without  action  on  indicators.  (4)  When  dissolved 
in  water  they  form  two  kinds  of  ions. 

Hfeat  of  neutralization.  If  neutralization  is  due  to  the  union  of 
hydrogen  ions  with  hydroxyl  ions,  and  nothing  more,  it  follows  that 
when  a  given  weight  of  water  is  formed  in  neutralization,  the  heat 
set  free  should  always  be  the  same,  no  matter  from  what  acid  and 
base  the  two  kinds  of  ions  have  been  supplied.  Careful  experiments 
have  shown  that  this  is  the  case,  provided  no  other  reactions  take 
place  at  the  same  time.  When  i8g.  of  water  are  formed  in  neu- 
tralization, 13,700  cal.  of  heat  are  set  free.  This  is  represented  in 
the  equations 

Na+,  OH-  +  H+,  Cl-  =  Na+,  Cl~  +  H2O  +  13,700  cal. 
K+,  OH-  +  H+,  NO-  =  K+,  NO-  +  H2O  +  13,700  cal. 
Ca+  +,  (OH)~-  +  H++,  SO—  =  Ca++,  SO"  +  2  H2O 
+  2  x  13,700  cal. 

Neutralization  a  quantitative  act.  Since  neutralization  is 
a  definite  chemical  act,  each  acid  will  require  a  perfectly 
definite  weight  of  each  base  for  its  neutralization.  For 


110     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


example,  a  given  weight  of  sulphuric  acid  will  always  require 
a  definite  weight  of  sodium  hydroxide,  in  accordance  with 
the  equation 

H2,  S04  +  2  Na,  OH  =  Na2,  SO4  +  2  H2O. 

Determination  of   the  ratio  in   neutralization.    The    quantities   of 
acid  and  base  required  in  neutralization  may  be  determined  in  the 
following  way.     Dilute  solutions  of   the    two 
substances  are  prepared, '  the    sulphuric    acid 
being  placed  in  one  of  the  burettes  (Fig.  34) 
and  the  sodium  hydroxide  in  the  other.    The 
levels  of  the  two  liquids  are  then  brought  to 
the  zero  marks  of  the  burettes  by  means  of  the 
stopcocks.    A  measured  volume  of  the  acid  is 
drawn  off  into  a  beaker,  a  few  drops  of  litmus 
solution   added,  and  the  sodium  hydroxide  is 
run  in  drop  by  drop  until  the  red  litmus  just 
turns  blue.    The  volume  of  the  sodium  hydrox- 
ide consumed  is  then  noted.    If  the  concentra- 
tions of  the  two  solutions  are  known,  it  is  easy 
to  calculate  what  weight  of  sodium  hydroxide 
is  required  to  neutralize  a  given  weight  of  sul- 
phuric acid.    By  evaporating  the  neutralized 
solution    to    dryness,    the    weight    of    the 
sodium  sulphate  formed  can  be  determined 

directly.  Experiment  shows  that  the  weights  are  always  in  accord- 
ance with  the  equation  in  the  preceding  paragraph. 

Extent  of  dissociation.  The  question  will  naturally  arise, 
When  an  acid,  base,  or  salt  dissolves  in  water,  do  all  the 
molecules  dissociate  into  ions,  or  only  a  part  of  them  ?  The 
experiments  by  which  this  question  can  be  answered  cannot 
be  described  here.  It  has  been  found,  however,  that  only  a 
fraction  of  the  molecules  dissociate.  The  percentage  which 
will  dissociate  in  a  given  case  depends  upon  several  con- 
ditions, the  chief  of  which  are :  ( I )  The  concentration  of 
the  solution.  In  concentrated  solutions  only  a  very  small 


ACIDS,  BASES,  AND  SALTS  1 1 1 

percentage  of  dissociation  occurs.  As  the  solution  is  diluted 
the  percentage  increases,  and  in  very  dilute  solutions  it  may 
be  very  large,  though  it  is  never  complete  in  any  ordinary 
solution.  (2)  The  nature  of  the  dissolved  compound.  At 
equal  concentrations  substances  differ  much  among  them- 
selves in  the  percentage  of  dissociation.  The  great  majority 
of  salts  are  about  equally  dissociated.  Acids  and  bases,  on 
the  contrary,  show  great  differences.  Some  are  freely  dis- 
sociated, while  others  are  dissociated  to  but  a  slight  extent. 
Strength  of  acids  and  bases.  Since  acid  and  basic  prop- 
erties are  due  to  hydrogen  and  hydroxyl  ions  respectively, 
the  acid  or  base  which  will  produce  the  greatest  percentage 
of  these  ions  at  a  given  concentration  must  be  regarded  as 
the  strongest  representative  of  its  class.  The  acids  and 
bases  described  in  the  foregoing  paragraphs  are  all  quite 
strong.  In  10%  solutions  they  are  dissociated  to  about  50%, 
and  this  is  also  approximately  the  extent  to  which  most 
salts  are  dissociated  at  this  same  concentration. 

Partial  neutralization.  I.  Basic  salts.  The  chemical  action  be- 
tween an  acid  and  a  base  is  not  always  as  complete  as  has  been 
represented  in  the  foregoing  paragraphs.  For  example,  if  the  base 
magnesium  hydroxide  (Mg(OH)2)  and  hydrochloric  acid  (HC1) 
are  brought  together  in  the  ratio  of  an  equal  number  of  molecules  of 
each,  there  will  be  only  half  enough  hydrogen  ions  for  the  hydroxyl 
ions  present.  » 

Mg,  (OH)2  +  H,  Cl  -  Mg,  OH,  Cl  +  H2O. 

Magnesium,  hydroxyl,  and  chlorine  ions  are  left  at  the  close  of  the 
reaction,  and  under  the  proper  conditions  unite  to  form  molecules 
of  the  compound  Mg(OH)Cl.  This  compound,  when  dissolved,  can 
form  hydroxyl  ions  and  therefore  possesses  basic  properties ;  it  can 
also  form  the  ions  of  a  salt  (Mg  and  Cl),  and  has  properties  char- 
acteristic of  salts.  Substances  of  this  kind  are  called  basic  salts. 

DEFINITION  :  A  basic  salt  is  a  substance  which  can  give  the  ions 
both  of  a  base  and  of  a  salt  when  dissolved  in  water. 


112     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

2.  Acid  salts.  In  a  similar  way,  when  sulphuric  acid  and  sodium 
hydroxide  are  brought  together  in  the  ratio  of  equal  numbers  of  the 
molecules  of  each,  it  is  possible  to  have  a  reaction  expressed  by  the 

Na,  OH  +  H2,  S04  =  Na,  H,  SO4  +  H2O. 

The  ions  remaining  after  all  the  hydroxyl  ions  have  been  used  up 
are  those  of  an  acid  (H)  and  those  of  a  salt  (Na  and  SO4).  These 
unite  to  form  the  substance  NaHSO4,  and  as  the  solution  becomes 
saturated  with  this  substance  through  evaporation,  it  separates  in  the 
form  of  crystals.  In  solution  this  substance  can  give  hydrogen  ions, 
and  therefore  possesses  acid  properties ;  it  can  also  give  the  ions 
characteristic  of  a  salt.  It  is  therefore  called  an  acid  salt. 

DEFINITION  :  An  acid  salt  is  one  which  can  give  the  ions  of  an 
acid  and  of  a  salt  when  in  solution. 

3.  Normal  salts.  Salts  which  are  the  products  of  complete  neu- 
tralization, such  as  Na2SO4,  and  which  in  .solution  can  give  neither 
hydrogen  nor  hydroxyl  ions,  but  only  the  ions  of  a  salt,  are  called 
normal  salts  to  distinguish  them  from  acid  and  basic  salts. 

Methods  of  expressing  reactions  between  compounds  in 
solution.  Chemical  equations  representing  reactions  between 
substances  in  solution  may  represent  the  details  of  the 
reaction,  or  they  may  simply  indicate  the  final  products 
formed.  In  the  latter  case  the  formation  of  ions  is  not 
indicated.  Thus,  if  we  wish  to  call  attention  to  the  details 
of  the  reaction  between  sodium  hydroxide  and  hydrochlo- 
ric acid  in  solution,  the  equation  is  written  as  follows : 

Na+,  OH-  +  H+,  Cl-  =  Na+,  Cl~  +  H2O. 

On  the  other  hand,  if  we  wish  simply  to  represent  the  final 
products  formed,  the  following  is  used. 

NaOH  +  HC1  =  NaCl  +  H2O. 

Both  of  these  methods  will  therefore  be  used : 

Radicals.  It  has  been  emphasized  that  the  hydroxyl 
group  (OH)  always  forms  the  anion  of  a  base,  while  the 


ACIDS,  BASES,  AND  SALTS         113 

group  NO3  forms  the  anion  of  nitric  acid  and  sodium 
nitrate ;  the  group  SO4,  the  anion  of  sulphuric  acid  and 
calcium  sulphate.  A  group  of  elements  which  in  this  way 
constitutes  a  part  of  a  molecule,  acting  as  a  unit  in  a 
chemical  change,  or  forming  ions  in  solution,  is  called  a 
radical.  Some  of  these  radicals  have  been  given  special 
names,  the  names  signifying  the  elements  present  in  the 
radical.  Thus  we  have  the  hydroxyl  radical  (OH)  and  the 
nitrate  radical  (NO3). 

DEFINITION  :  A  radical  is  a  group  of  elements  forming 
part  of  a  molecule,  and  acting  as  a  unit  in  chemical  reactions. 

Names  of  acids,  bases,  and  salts.  Since  acids,  bases, 
and  salts  are  so  intimately  related  to  each  other,  it  is  very 
advantageous  to  give  names  to  the  three  classes  in  accord- 
ance with  some  fixed  system.  The  system  universally 
adopted  is  as  follows : 

Naming  of  bases.  All  bases  are  called  hydroxides.  They, 
are  distinguished  from  each  other  by  prefixing  the  name  of 
the  element  which  is  in  combination  with  the  hydroxyl 
group.  Examples :  sodium  hydroxide  (NaOH) ;  calcium 
hydroxide  (Ca(OH)2) ;  copper  hydroxide  (Cu(OH)2). 

Naming  of  acids.  The  method  of  naming  acids  depends 
upon  whether  the  acid  consists  of  two  elements  or  three. 

1.  Binary  acids.    Acids  containing  only  one  element  in 
addition    to  hydrogen  are    called  binary  acids.    They  are 
given  names  consisting  of  the  prefix  hydro-,  the  name  of 
the  second  element  present,  and  the  termination  -ic.    Exam- 
ples :  hydrochloric  acid  (HC1)  ;  hydrosulphuric  acid  (H2S). 

2.  Ternary  acids.    In  addition  to  the  two  elements  pres- 
ent in  binary  acids,  the  great  majority  of  acids  also  con- 
tain oxygen.    They  therefore  consist  of  three  elements  and 


114     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

are  called  ternary  acids.  It  usually  happens  that  the  same 
three  elements  can  unite  in  different  proportions  to  make 
several  different  acids.  The  most  familiar  one  of  these  is 
given  a  name  ending  in  the  suffix  -ic,  while  the  one  with 
less  oxygen  is  given  a  similar  name,  but  ending  in  the  suffix 
-ous.  Examples  :  nitric  acid  (HNO3) ;  nitrous  acid  (HNO2). 
In  cases  where  more  than  two  acids  are  known,  use  is  made 
of  prefixes  in  addition  to  the  two  suffixes  -ic  and  -ous.  Thus 
the  prefix  per-  signifies  an  acid  still  richer  in  oxygen ;  the 
prefix  hypo-  signifies  one  with  less  oxygen. 

Naming  of  salts.  A  salt  derived  from  a  binary  acid  is 
given  a  name  consisting  of  the  names  of  the  two  elements 
composing  it,  with  the  termination  -ide.  Example  :  sodium 
chloride  (NaCl).  All  other  binary  compounds  are  named  in 
the  same  way. 

A  salt  of  a  ternary  acid  is  named  in  accordance  with  the 
acid  from  which  it  is  derived.  A  ternary  acid  with  the 
termination  -ic  gives  a  salt  with  the  name  ending  in  -ate, 
while  an  acid  with  termination  -ous  gives  a  salt  with  the 
name  ending  in  -ite.  The  following  table  will  make  the  appli- 
cation of  these  principles  clear : 

ACIDS  SYMBOL  SALTS  SYMBOL 

Hydrochloric  HC1  Sodium  chloride  NaCl 

Hypochlorous  HC1O  Sodium  hypochlorite  NaCIO 

Chlorous  HC1O2  Sodium  chlorite  NaClO2 

Chloric  HC1O3  Sodium  chlorate  NaClO3 

Perchloric  HC1O4  Sodium  perchlorate  NaClO4 


ACIDS,  BASES,  AND  SALTS  115 


EXERCISES 

1.  25  cc.  of  a  solution  containing  40  g.  of  sodium  hydroxide  per 
liter  was  found   to   neutralize  25  cc.  of  a  solution  of  hydrochloric 
acid.    What  was  the  strength  of  the  acid  solution  ?      H 

2.  After  neutralizing  a  solution  of  sodium  hydroxide  with  nitric 
acid,  there   remained   after   evaporation    i  oo  g.  of   sodium   nitrate. 
How  much  of  each  substance  had  been  used  ?  4^[ 

3.  A  solution  contains  18  g.  of  hydrochloric  acid  per  100  cc.     It 
required  25  cc.  of  this  solution  to  neutralize  30  cc.  of  a  solution  of 
sodium  hydroxide.    What  was  the  strength  of  the  sodium  hydroxide 
solution  in  parts  per  hundred? 

4.  When  perfectly  dry  sulphuric  acid  is  treated  with  perfectly 
dry  sodium  hydroxide,  no  chemical  change  takes  place.    Explain.    *"**" 

5.  When  cold,  concentrated  sulphuric  acid  is  added  to  zinc,  no 
change  takes  place.    Recall  the  action  of  dilute  sulphuric  acid  on 
the  same  metal.     How  do  you  account  for  the  difference  ? 

6.  A  solution  of  hydrochloric  acid  in  benzene  does  not  conduct 
the  electric  current.    When  this  solution  is  treated  with  zinc,  will  « 
hydrogen  be  evolved  ?    Explain. 

7.  (a)  Write  equation  for  preparation  of  hydrogen  from  zinc  and 
dilute   sulphuric    acid.     (£)  Rewrite    the    same    equation   from   the     ^  ' 
standpoint  of   the  theory  of  electrolytic  dissociation.    (V)  Subtract 

the  common  SO4  ion  from  both  members  of  the  equation.  (*/)  From 
the  resulting  equation,  explain  in  what  the  preparation  of  hydrogen 
consists  when  examined  from  the  standpoint  of  this  theory. 

8.  In  the  same  manner  as  in  the  preceding  exercise,  explain  in 
what  the  action  of  sodium  on  water  to  give  hydrogen  consists. 

•frvl^s.* 


£& 

vi      -^^°^ 

r  f  f 

- 


CHAPTER  XI 
VALENCE 

Definition  of  valence.  A  study  of  the  formulas  of  various 
binary  compounds  shows  that  the  elements  differ  between 
themselves  in  the  number  of  atoms  of  other  elements  which 
they  are  able  to  hold  in  combination.  This  is  illustrated  in 
the  formulas 

HC1,  H2O,  H3N,  H4C. 

(hydrochloric  acid)  (water)  (ammonia)  (marsh  gas) 

It  will  be  noticed  that  while  one  atom  of  chlorine  combines 
with  one  atom  of  hydrogen,  an  atom  of  oxygen  combines 
with  two,  an  atom  of  nitrogen  with  three,  one  of  car- 
bon with  four.  The  number  which  expresses  this  com- 
bining ratio  between  atoms  is  a  definite  property  of  each 
element  and  is  called  its  valence. 

DEFINITION  : .  The  valence  of  an  element  is  that  property 
which  determines  the  number  of  the  atoms  of  another  ele- 
ment which  its  atom  can  Jiold  in  combination. 

Valence  a  numerical  property.  Valence  is  therefore 
merely  a  numerical  relation  and  does  not  convey  any  in- 
formation in  regard  to  the  intensity  of  the  affinity  between 
atoms.  Judging  by  the  heat  liberated  in  their  union,  oxygen 
has  a  far  stronger  affinity  for  hydrogen  than  does  nitrogen, 
but  an  atom  of  oxygen  can  combine  with  two  atoms  only 
of  hydrogen,  while  an  atom  of  nitrogen  can  combine  with 
three. 

116 


VALENCE  117 

Measure  of  valence.  In  expressing  the  valence  of  an 
element  we  must  select  some  standard  for  comparison,  just 
as  in  the  measurement  of  any  other  numerical  quantity.  It 
has  been  found  that  an  atom  of  hydrogen  is  never  able  to 
hold  in  combination  more  than  one  atom  of  any  other  ele- 
ment. Hydrogen  is  therefore  taken  as  the  standard,  and 
other  elements  are  compared  with  it  in  determining  their 
valence.  A  number  of  other  elements  are  like  hydrogen  in 
being  able  to  combine  with  at  most  one  atom  of  other  ele- 
ments, and  such  elements  are  called  univalent.  Among 
these  are  chlorine,  iodine,  and  sodium.  Elements  such  as 
oxygen,  calcium,  and  zinc,  which  can  combine  with  two 
atoms  of  hydrogen  or  other  univalent  elements,  are  said  to 
be  divalent.  Similarly,  we  have  trivalent,  tetravalent,  pen- 
tavalent  elements.  None  have  a  valence  of  more  than  8. 

Indirect  measure  of  valence.  Many  elements,  especially 
among  the  metals,  do  not  readily  form  compounds  with 
hydrogen,  and  their  valence  is  not  easy  to  determine  by 
direct  comparison  with  the  standard  element.  These  ele- 
ments, however,  combine  with  other  univalent  elements, 
such  as  chlorine,  and  their  valence  can  be  determined 
from  the  compounds  so  formed. 

Variable  valence.  Many  elements  are  able  to  exert  differ- 
ent valences  under  differing  circumstances.  Thus  we  have 
the  compounds  Cu2O  and  CuO,  CO  and  CO2,  FeCl2  and 
FeCl3.  It  is  not  always  possible  to  assign  a  fixed  valence 
to  an  element.  Nevertheless  each  element  tends  to  exert 
some  normal  valence,  and  the  compounds  in  which  it  has  a 
valence  different  from  this  are  apt  to  be  unstable  and  easily 
changed  into  compounds  in  which  the  valence  of  the  ele- 
ment is  normal.  The  valences  of  the  various  elements  will 
become  familiar  as  the  elements  are  studied  in  detail. 


118     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Valence  and  combining  ratios.  When  elements  combine 
to  form  compounds,  the  ratio  in  which  they  combine  will  be 
determined  by  their  valences.  In  those  compounds  which 
consist  of  two  elements  directly  combined,  the  union  is 
between  such  numbers  of  the  two  atoms  as  have  equal 
valences.  Elements  of  the  same  valence  will  therefore 
combine  atom  for  atom.  Designating  the  valence  of  the 
atoms  by  Roman  numerals  placed  above  their  symbols,  we 

have  the  formulas 

ii  n  ii  mm  iv iv 

HC1,  ZnO,  BN,  CSi. 

A  divalent  element,  on  the  other  hand,  will  combine  with  two 

ii   ii  n   ii 

atoms  of  a  univalent  element.    Thus  we  have  ZnCl2  and  H2O 

(the   numerals  above  each   symbol  representing  the  sum 
of  the  valences  of  the  atoms  of  the  element  present).    A 

trivalent  atom  will  combine  with  three  atoms  of  a  univalent 

in  m 
element,  as  in  the  compound  H3N. 

If  a  trivalent  element  combines  with  a  divalent  element, 
the  union  will  be  between  two  atoms  of  the  trivalent  element 
and  three  of  the  divalent  element,  since  these  numbers  are 
the  smallest  which  have  equal  valences.  Thus  the  oxide 
of  the  trivalent  metal  aluminium  has  the  formula  A12O3. 
Finally  one  atom  of  a  tetravalent  element  such  as  carbon 
will  combine  witR  four  atoms  of  a  univalent  element,  as  in 
the  compound  CH4,  or  with  two  atoms  of  a  divalent  element, 
as  in  the  compound  CO2. 

We  have  no  knowledge  as  to  why  elements  differ  in  their 
combining  power,  and  there  is  no  way  to  determine  their 
valences  save  by  experiment. 

Valence  and  the  structure  of  compounds.  Compounds  will  be  met 
from  time  to  time  which  are  apparent  exceptions  to  the  general  state- 
ments just  made  in  regard  to  valence.  Thus,  from  the  formula  for 


VALENCE  119 

hydrogen  dioxide  (H2O2),  it  might  be  supposed  that  the  oxygen  is 
univalent  ;  yet  it  is  certainly  divalent  in  water  (H2O).  That  it  may 
also  be  divalent  in  H2O2  may  be  made  clear  as  follows:  The  unit 
valence  of  each  element  may  be  represented  graphically  by  a  line 
attached  to  its  symbol.  Univalent  hydrogen  and  divalent  oxygen 
will  then  have  the  symbols  H—  and  —  O  —  .  When  atoms  combine, 
each  unit  valence  of  one  atom  combines  with  a  unit  valence  of  another 
atom.  Thus  the  composition  of  water  may  be  expressed  by  the  for- 
mula H  —  O  —  H,  which  is  meant  to  show  that  each  of  the  unit  valences 
of  oxygen  is  satisfied  with  the  unit  valence  of  a  single  hydrogen 
atom. 

The  chemical  conduct  of  hydrogen  dioxide  leads  to  the  conclusion 
that  the  two  oxygen  atoms  of  its  molecule  are  in  direct  combination 
with  each  other,  and  in  addition  each  is  in  combination  with  a  hydro- 
gen atom.  This  may  be  expressed  by  the  formula  H  —  O  —  O  —  H. 
The  oxygen  in  the  compound  .is  therefore  divalent,  just  as  it  is  in 
water.  It  will  thus  be  seen  that  the  structure  of  a  compound  must 
be  known  before  the  valences  of  the  atoms  making  up  the  compound 
can  be  definitely  decided  upon. 

Such  formulas  as  H  —  O  —  H  and  H  —  O  —  O  —  H  are  known  as 
structural  formulas,  because  they  are  intended  to  show  what  is  known 
in  regard  to  the  arrangement  of  the  atoms  in  the  molecules. 

Valence  and  the  replacing  power  of  atoms.  Just  as  ele- 
ments having  the  same  valence  combine  with  each  other 
atom  for  atom,  so  if  they  replace  each  other  in  a  chemical 
reaction  they  will  do  so  in  the  same  ratio.  This  is  seen  in 
the  following  equations,  in  which  a  univalent  hydrogen  atom 
is  replaced  by  a  univalent  sodium  atom  : 

NaOH  -f  HC1  =  NaCl  +  H2O. 
2  NaOH  +  H2SO4  -  Na2SO4  +  2  H2O. 


Similarly,  one  atom  of  divalent  calcium  will  replace  two 
atoms  of  univalent  hydrogen  or  one  of  divalent  zinc  : 

Ca(OH)2  +  2  HC1  =  CaQ3  +  2  H2O. 
CaCl2  +  ZnSO4  =  CaSO4  +  ZnCl2. 


1 20     AN  ELEMENTARY  STUDY  OF  CHEMISTRY  . 

In  like  manner,  one  atom  of  a  trivalent  element  will  replace 
three  of  a  univalent  element,  or  two  atoms  will  replace  three 
atoms  of  a  divalent  element. 

Valence  and  its  applications  to  formulas  of  salts.  While  the  true 
nature  of  valence  is  not  understood  and  many  questions  connected 
with  the  subject  remain  unanswered,  yet  many  of  the  main  facts  are 
of  much  help  to  the  student.  Thus  the  formula  of  a  salt  differs  from 
that  of  the  acid  from  which  it  is  derived  in  that  the  hydrogen  of  the 
acid  has  been  replaced  by  a  metal.  If,  then,  it  is  known  that  a  given 
metal  forms  a  normal  salt  with  a  certain  acid,  the  formula  of  the  salt 
can  at  once  be  determined  if  the  valence  of  the  metal  is  known.  Since 
sodium  is  univalent,  the  sodium  salts  of  the  acids  HC1  and  H2SO4 
will  be  respectively  NaCl  and  Na2SO4.  One  atom  of  divalent 
zinc  will  replace  2  hydrogen  atoms,  so  that  the  corresponding  zinc 
salts  will  be  ZnCl2  and  ZnSO4. 

The  formula  for  aluminium  sulphate  is  somewhat  more  difficult  to 
determine.  Aluminium  is  trivalent,  and  the  simplest  ratio  in  which 
the  aluminium  atom  can  replace  the  hydrogen  in  sulphuric  acid  is 
2  atoms  of  aluminium  (6  valences)  to  3  molecules  of  sulphuric  acid 
(6  hydrogen  atoms).  The  formula  of  the  sulphate  will  then  be 
A12(S04)3. 

Valence  and  its  application  to  equation  writing.  It  will  be  readily 
seen  that  a  knowledge  of  valence  is  also  of  very  great  assistance  in 
writing  the  equations  for  reactions  of  double  decomposition.  Thus, 
in  the  general  reaction  between  an  acid  and  a  base,  the  essential 
action  is  between  the  univalent  hydrogen  ion  and  the  univalent 
hydroxyl  ion.  The  base  and  the  acid  must  always  be  taken  in  such 
proportions  as  to  secure  an  equal  number  of  each  of  these  ions.  Thus, 
in  the  reaction  between  ferric  hydroxide  (Fe(OH)3)  and  sulphuric  acid 
(H2SO4),  it  will  be  necessary  to  take  2  molecules  of  the  former  and  3 
of  the  latter  in  order  to  have  an  equal  number  of  the  two  ions,  namely, 
6.  The  equation  will  then  be 

2  Fe(OH)3  +  3  H2S04  =  Fe2(SO4)3  +  6  H2O. 

Under  certain  conditions  the  salts  A12(SO4)3  and  CaCl2  undergo 
double  decomposition,  the  two  metals,  aluminium  and  calcium,  ex- 
changing places.  The  simplest  ratio  of  exchange  in  this  case  is  2 
atoms  of  aluminium  (6  valences)  and  3  atoms  of  calcium  (6  valences). 


VALENCE  1 2 1 

The  reaction  will  therefore  take  place  between  i  molecule  of  A12(SO4)3 
and  3  of  CaCl2,  and  the  equation  is  as  follows: 

A12(SO4)3  +  3  CaCl2  =  3  CaSO4  +  2  A1C13. 


EXERCISES 

1.  Sodium,  calcium,  and  aluminium  have  valences  of  1,2,  and  3 
respectively ;  write  the  formulas  of  their  chlorides,  sulphates,  and 
phosphates  (phosphoric  acid  =  H3PO4),  on  the  supposition  that  they 
form  salts  having  the  normal  composition. 

2.  Iron  forms  one  series  of  salts  in  which  it  has  a  valence  of  2, 
and  another  series  in  which  it  has  a  valence  of  3  ;  write  the  formulas' 
for  the  two  chlorides  of  iron,  also  for  the  two  sulphates,  on  the  suppo- 
sition that  these  have  the  normal  composition. 

3.  Write  the  equation  representing  the  neutralization  of  each  of 
the  following  bases  by  each  of  the  acids  whose  formulas  are  given  : 

NaOH  HC1 

Ba(OH)2  H2S04 

A1(OH)3  H3P04 

4.  Silver  acts  as  a  univalent  element  and  calcium  as  a  divalent 
element  in  the  formation  of  their  respective  nitrates  and  chlorides. 
(a)    Write    the    formula    for   silver   nitrate  ;    for   calcium   chloride. 
(£)  When  solutions  of  these  two  salts  are  mixed,  the  two  metals, 
silver  and   calcium,  exchange   places ;    write   the   equation  for  the 
reaction. 

5.  Antimony  acts  as  a  trivalent  element  in  the  formation  of  a 
chloride.    (#)  What  is  the  formula  for  antimony  chloride  ?  (£)  When 
hydrosulphuric  acid  (H2S)  is  passed  into  a  solution  of  this  chloride 
the  hydrogen  and  antimony  exchange  places ;  write  the  equation  for 
the  reaction. 

6.  Lead  has  a  valence  of  2  and  iron  of  3  in  the  compounds  known 
respectively  as  lead  nitrate  and  ferric  sulphate,    (a)  Write  the  for- 
mulas for  these  two  compounds.    (6)  When  their  solutions  are  mixed 
the  two  metals  exchange  places ;  write  the  equation  for  the  reaction. 


CHAPTER   XII 
COMPOUNDS  OF  NITROGEN 

Occurrence.  As  has  been  stated  in  a  former  chapter, 
nitrogen  constitutes  a  large  fraction  of  the  atmosphere. 
'The  compounds  of  nitrogen,  however,  cannot  readily  be 
obtained  from  this  source,  since  at  any  ordinary  temper- 
ature nitrogen  is  able  to  combine  directly  with  very  few  of 
the  elements. 

In  certain  forms  of  combination  nitrogen  occurs  in  the 
soil  from  which  it  is  taken  up  by  plants  and  built  into 
complex  substances  composed  chiefly  of  carbon,  hydrogen, 
oxygen,  and  nitrogen.  Animals  feeding  on  these  plants  as- 
similate the  nitrogenous  matter,  so  that  this  element  is  an 
essential  constituent  of  both  plants  and*animals. 

Decomposition  of  organic  matter  by  bacteria.  When  liv- 
ing matter  dies  and  undergoes  decay  complicated  chemical 
reactions  take  place,  one  result  of  which  is  that  the  nitro- 
gen of  the  organic  matter  is  set  fr.ee  either  as  the  element 
nitrogen,  or  in  the  form  of  simple,  compounds,  such  as 
ammonia  (NH3)  or  oxides  of  nitrogen.  Experiment  has 
shown  that  all  such  processes  of  decay  are  due  to  the 
action  of  different  kinds  of  bacteria,  each  particular  kind 
effecting  a  different  change. 

Decomposition  of  organic  matter  by  heat.  When  organic 
matter  is  strongly  heated  decomposition  into  simpler  sub- 
stances takes  place  in  much  the  same  way  as  in  the  case  of 
bacterial  decomposition.  Coal  is  a  complex  substance  of 

122 


COMPOUNDS  OF  NITROGEN 


I23 


vegetable  origin,  consisting  largely  of  carbon,  but  also  con- 
taining hydrogen,  oxygen,  and  nitrogen.  When  this  is 
heated  in  a  closed  vessel  so  that  air  is  excluded,  about 
one  seventh  of  the  nitrogen  is  converted  into  ammonia, 
and  this  is  the  chief  source  from  which  ammonia  and  its 
compounds  are  obtained. 

COMPOUNDS  OF  NITROGEN  WITH  HYDROGEN 

Ammonia  (NH8).  Several  compounds  consisting  exclu- 
sively of  nitrogen  and  hydrogen  are  known,  but  only  one, 
ammonia,  need  be  considered  here. 

Preparation  of  ammonia.  Ammonia  is  prepared  in  the 
laboratory  by  a  different  method  from  the  oner  which  is 
used  commercially. 

i.  Laboratory  method.  In  the  laboratory  ammonia  is 
prepared  from  ammonium  chloride,  a  compound  having  the 
formula  NH4C1,  and  obtained  in  the  manufacture  of  coal 
gas.  As  will  be  sr^pwn  later  in  the  chapter,  the  group  NH4 
in  this  compound  acts  as  a  univalent  radical  and  is  known  as 
ammonium.  When  ammonium  chloride  is  warmed  with  so- 
dium hydroxide,  the  ammonium  and  sodium  change  places, 
the  reaction  being  expressed  in  the  following  equation. 

NH4C1  +  NaOH  =  NaCl  +  NH4OH. 

The  ammonium  hydroxide  (NH4OH)  so  formed  is  unstable, 
and  breaks  down  into  water  and  ammonia. 

NH4OH  =  NH3  +  H2O. 

Calcium  hydroxide  (Ca(OH)2)  is  frequently  used  in  place  of 
the  more  expensive  sodium  hydroxide,  the  equations  being 

2  NH4C1  +  Ca(OH)2  =  CaCl2  +  2  NH4OH, 
2  NH4OH  =  2  H2O  +  2  NH3. 


124     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


In  the  preparation,  the  ammonium  chloride  and  calcium  hydroxide 
are  mixed  together  and  placed  in  a  flask  arranged  as  shown  in  Fig.  35. 

The  mixture  is  gently  warmed, 
when  ammonia  is  evolved  as  a 
gas  and  is  collected  by  displace- 
ment of  air. 

2.  Commercial  method. 
Nearly  all  the  ammonia  of 
commerce  comes  from  the 
gasworks.  Ordinary  illumi- 
nating gas  is  made  by  dis- 
tilling coal,  as  will  be 
explained  later,  and  among 
the  products  of  this 

-  distillation  a  solution 
of  ammonia  in  water 


FIG.  35 


is  obtained.  This  solution,  known  as  gas  liquor,  contains 
not  only  ammonia  but  other  soluble  substances.  Most  of 
these  combine  chemically  with  lime,  while  ammonia  does 
not ;  if  then  lime  is  added  to  the  gas  liquor  and  the  liquor 
is  heated,  the  ammonia  is  driven  out  from  the  mixture. 
It  may  be  dissolved  again  in  pure,  cold  water,  forming  aqua 
ammonia,  or  the  ammonia  water  of  commerce. 

Preparation  from  hydrogen  and  nitrogen.  When  electric  sparks  are 
passed  for  some  time  through  a  mixture  of  hydrogen  and  nitrogen,  a 
small  percentage  of  the  two  elements  in  the  mixture  is  changed  into 
ammonia.  The  action  soon  ceases,  however,  for  the  reason  that 
ammonia  is  decomposed  by  the  electric  discharge.  The  reaction 
expressed  in  the  equation 

N  +  sH  =  NH3 

can  therefore  go  in  either  direction  depending  upon  the  relative  quan- 
tities of  the  substances  present.  This  recalls  the  similar  change  from 
oxygen  into  ozone,  which  soon  ceases  because  the  ozone  is  in  turn 
decomposed  into  oxygen. 


COMPOUNDS  OF  NITROGEN  125 

Physical  properties.  Under  ordinary  conditions  ammonia 
is  a  gas  whose  density  is  0.59.  It  is  therefore  little  more 
than  half  as  heavy  as  air.  It  is  easily  condensed  into  a 
colorless  liquid,  and  can  now  be  purchased  in  liquid  form 
in  steel  cylinders.  The  gas  is  colorless  and  has  a  strong, 
suffocating  odor.  It  is  extremely  soluble  in  water,  i  1.  of 
water  at  o°  and  760  mm.  pressure  dissolving  1 148  1.  of  the 
gas.  In  dissolving  this  large  volume  of  gas  the  water  ex- 
pands considerably,  so  that  the  density  of  the  solution  is 
less  than  that  of  water,  the  strongest  solutions  having  a 
density  of  0.88. 

Chemical  properties.  Ammonia  will  not  support  com- 
bustion, nor  will  it  burn  under  ordinary  conditions.  In  an 
atmosphere  of  oxygen  it  burns  with  a  feeble,  yellowish 
flame.  When  quite  dry  it  is  not  a  very  active  substance, 
but  when  moist  it  combines  with  a  great  many  substances, 
particularly  with  acids. 

Uses.  It  has  been  stated  that  ammonia  can  be  con- 
densed to  a  liquid  by  the  application  of  pressure.  If  the 
pressure  is  removed  from  the  liquid  so  obtained,  it  rapidly 
passes  again  into  the  gaseous  state  and  in  so  doing  absorbs 
a  large  amount  of  heat.  Advantage  is  taken  of  this  fact 
in  the  preparation  of  artificial  ice.  Large  quantities  of 
ammonia  are  also  used  in  the  preparation  of  ammonium 
compounds. 

The  manufacture  of  artificial  ice.  Fig.  36  illustrates  the  method  of 
preparing  artificial  ice.  The  ammonia  gas  is  liquefied  in  the  pipes  X 
by  means  of  the  pump  Y.  The  heat  generated  is  absorbed  by  water 
flowing  over  the  pipes.  The  pipes  lead  into  a  large  brine  tank,  a 
cross  section  of  which  is  shown  in  the  figure.  Into  the  brine  (con- 
centrated solution  of  common  salt)  contained  in  this  tank  are  dipped 
the  vessels  A,  B,  C,  filled  with  pure  water.  The  pressure  is  removed 
from  the  liquid  ammonia  as  it  passes  into  the  pipes  immersed  in  the 


126     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


Water 
Ammonia  ^  Condenser 


brine,  and  the  heat  absorbed  by  the  rapid  evaporation  of  the  liquid 
lowers  the  temperature  of  the  brine  below  zero.    The  water  in  A,  £,  C 

is  thereby  frozen 
into  cakes  of  ice. 
The  gaseous  am- 
monia resulting 
from  the  evapora- 
tion of  the  liquid 
ammonia  is  again 
condensed,  so  that 
the  process  is  con- 
tinuous. 


Compressor 
Pump  Y 


Brine  Tank  (Section) 
A,  B,C,  Ice  Tanks    D,  Brine 

FIG.  36 


Ammonium 
hydroxide 

(NH4OH).  The 


solution  of  ammonia  in  water  is  found  to  have  strong  basic 
properties  and  therefore  contains  hydroxyl  ions.  It  turns 
red  litmus  blue ;  it  has  a  soapy  feel ;  it  neutralizes  acids, 
forming  salts  with  them.  It  seems  probable,  therefdreVthat 
when  ammonia  dissolves  in  water  it  combines  chemically 
with  it  according  to  the  equation 

NH3  +  H2O  =  NH4OH, 

and  that  it  is  the  substance  NH4OH,  called  ammonium 
hydroxide,  which  has  the  basic  properties,  dissociating  into 
the  ions  NH4  and  OH.  Ammonium  hydroxide  has  never 
been  obtained  in  a  pure  state.  At  every  attempt  to  isolate 
it  the  substance  breaks  up  into  water  and  ammonia,  — 

NH4OH  =  NH3  +  H2O. 

The  ammonium  radical.  The  radical  NH4  plays  the  part 
of  a  metal  in  many  chemical  reactions  and  is  called  ammo- 
nium. The  ending  -ium  is  given  to  the  name  to  indicate 
the  metallic  properties  of  the  substance,  since  the  names 


COMPOUNDS  OF  NITROGEN  127 

of  the  metals  in  general  have  that  ending.  The  salts  formed 
by  the  action  of  the  base  ammonium  hydroxide  on  acids 
are  called  ammonium  salts.  Thus,  with  hydrochloric  acid, 
ammonium  chloride  is  formed  in  accordance  with  the 
equatic-n  Nuf>u  +  HQ  =  NH^Q  +  ^ 

Similarly,  with  nitric  acid,  ammonium  nitrate  (NH4NO3)  is 
formed,  and  with  sulphuric  acid,  ammonium  sulphate 
((NH4)2S04). 

It  will  be  noticed  that  in  the  neutralization  of  ammo- 
nium hydroxide  by  acids  the  group  NH4  replaces  one  hydro- 
gen atom  of  the  acid,  just  as  sodium  does.  The  group 
therefore  acts  as  a  univalent  metal. 

Combination  of  nitrogen  with  hydrogen  by  volume. 
Under  suitable  conditions  ammonia  can  be  decomposed 
into  nitrogen  and  hydrogen  by  passing  electric  sparks 
through  the  gas.  Accurate  measurement  has  shown  that 
when  ammonia  is  decomposed,  two  volumes  of  the  gas 
yield  one  volume  of  nitrogen  and  three  volumes  of  hydro- 
gen. Consequently,  if  the  two  elements  were  to  combine 
directly,  one  volume  of  nitrogen  would  combine  with  three 
volumes  of  hydrogen  to  form  two  volumes  of  ammonia. 
Here,  as  in  the  formation  of  steam  from  hydrogen  and 
oxygen,  small  whole  numbers  serve  to  indicate  the  rela- 
tion between  the  volumes  of  combining  gases  and  that  of 
the  gaseous  product. 

COMPOUNDS  OF  NITROGEN  WITH  OXYGEN  AND  HYDROGEN 

In  addition  to  ammonium  hydroxide,  nitrogen  forms 
several  compounds  with  hydrogen  and  oxygen,  of  which 
nitric  acid  (HNO3)  and  nitrous  acid  (HNO2)  are  the  most 
familiar. 


128     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


Nitric  acid  (HNO3).  Nitric  acid  is  not  found  to  any 
extent  in  nature,  but  some  of  its  salts,  especially  sodium 
nitrate  (NaNO3)  and  potassium  nitrate  (KNO3)  are  found 
in  large  quantities.  From  these  salts 
nitric  acid  can  be  obtained. 

Preparation  of  nitric  acid.    When 
sodium  nitrate  is  treated  with  concen- 
trated cold  sulphuric  acid,  no  chemical 
action  seems  to  take  place.   If,  how- 
ever, the  mixture  is  heated  in  a  retort, 
nitric  acid  is  given  off 
as  a  vapor  and  may  be 
easily  condensed   to  a 
liquid   by  passing   the 
vapor  into  a  tube  sur- 
rounded by  cold  water, 


FIG. 37 


as  shown  in  Fig.  37.  An  examination  of  the  liquid  left  in  the 
retort  shows  that  it  contains  sodium  acid  sulphate  (NaHSO4), 
so  that  the  reaction  may  be  represented  by  the  equation 

NaN03  +  H2S04  =  NaHSO4  +  HNO3. 

If  a  smaller  quantity  of  sulphuric  acid  is  taken  and  the  mixture 
is  heated  to  a  higji  temperature,  normal  sodium  sulphate  is  formed  : 

2  NaNO8  +  H2SO4  =  Na2SO4  +  2  HNO3. 

In  this  case,  however,  the  higher  temperature  required  decomposes 
a  part  of  the  nitric  acid. 

The  commercial  preparation  of  nitric  acid.  Fig.  38  illustrates  a 
form  of  apparatus  used  in  the  preparation  of  nitric  acid  on  a  large 
scale.  Sodium  nitrate  and  sulphuric  acid  are  heated  in  the  iron 
retort  A.  The  resulting  acid  vapors  pass  in  the  direction  indicated 
by  the  arrows,  and  are  condensed  in  the  glass  tubes  B,  which  are 
covered  with  cloth  kept  cool  by  streams  of  water.  These  tubes 
are  inclined  so  that  the  liquid  resulting  from  the  condensation  of  the 
vapors  runs  back  into  C  and  is  drawn  off  into  large  vessels  (Z?). 


COMPOUNDS  OF  NITROGEN 


129 


Physical  properties  of  nitric  acid.  Pure  nitric  acid  is  a 
colorless  liquid,  which  boils  at  about  86°  and  has  a  density 
of  1.56.  The  concentrated  acid  of  commerce  contains  about 
68%  of  the  acid,  the  remainder  being  water.  Such  a  mix- 
ture has  a  density  of  1.4.  The  concentrated  acid  fumes 
somewhat  in  moist 
air,  and  has  a  sharp 
choking  odor. 

Chemical    proper- 
ties.   The  most  im- 
portant chemical 
properties    of 
nitric  acid  are  the 
following. 

i.  Acid  prop- 
erties.   As  the 
indicates. 


-NaNOj 
H2SO4 


name  indicates, 
this  substance  is  an  acid,  and  has  all  the  properties  of  that 
class  of  substances.  It  changes  blue  litmus  red  and  has  a 
sour  taste  in  dilute  solutions.  It  forms  hydrogen  ions  in 
solution  and  neutralizes  bases  forming  salts.  It  also  acts 
upon  the  oxides  of  most  metals,  forming  a  salt  and  water. 
It  is  one  of  the  strongest  acids. 

2.  Decomposition  on  heating.  When  boiled,  or  exposed 
for  some  time  to  sunlight,  it  suffers  a  partial  decomposition 
according  to  the  equation 

2  HNO3  =  H2O  +  2  NO2  +  O. 

The  substance  NO2,  called  nitrogen  peroxide,  is  a  brown- 
ish gas,  which  is  readily  soluble  in  water  and  in  nitric 
acid.  It  therefore  dissolves  in  the  undecomposed  acid,  and 
imparts  a  yellowish  or  reddish  color  to  it.  Concentrated 


130     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

nitric  acid  highly  charged  with  this  substance  is  called 
fuming  nitric  acid. 

3.  Oxidizing  action.   According  to  its  formula,  nitric  acid 
contains  a  large  percentage  of  oxygen,  and  the  reaction  just 
mentioned  shows  that  the  compound  is  not  a  very  stable 
one,    easily    undergoing   decomposition.    These    properties 
should  make  it  a  good  oxidizing  agent,  and  we  find  that  this 
is  the  case.    Under  ordinary  circumstances,   when  acting 
as  an  oxidizing  agent,  it  is  decomposed  according  to  the 

equation 

2  HNO3=  H2O  +  2  NO  +  3  O. 

The  oxygen  is  taken  up  by  the  substance  oxidized,  and  not 
set  free,  as  is  indicated  in  the  equation.  Thus,  if  carbon  is 
oxidized  by  nitric  acid,  the  oxygen  combines  with  carbon, 
forming  carbon  dioxide  (CO2) : 

C  +  2  O  -  C02. 

4.  Action  on  metals.    We  have  seen  that  when  an  acid 
acts  upon  a  metal  hydrogen  is  set  free.    Accordingly,  when 
nitric  acid  acts  upon  a  metal,  such  as  copper,  we  should 
expect  the  reaction  to  take  place  which  is  expressed  in  the 

equation 

Cu  +  2  HNO3  -  Cu(NO3)2  4-  2  H. 

This  reaction  does  take  place,  but  the  hydrogen  set  free  is 
immediately  oxidized  to  water  by  another  portion  of  the 
nitric  acid  according  to  the  equation 

HNO3  +  3  H  =  2  H2O  H-  NO. 

As  these  two  equations  are  written,  two  atoms  of  hydrogen 
are  given  off  in  the  first  equation,  while  three  are  used  up 
in  the  second.  In  order  that  the  hydrogen  may  be  equal  in 


COMPOUNDS  OF  NITROGEN  131 

the  two  equations,  we  must  multiply  the  first  by  3  and  the 
second  by  2.  We  shall  then  have 

3  Cu  +  6  HN03  =  3  Cu(N03)2  +  6  H, 
2  HNO3  +  6  H  =  4  H2O  +  2  NO. 

The  two  equations  may  now  be  combined  into  one  by  add- 
ing the  quantities  on  each  side  of  the  equality  sign,  can- 
celing the  hydrogen  which  is  given  off  in  the  one  reaction 
and  used  up  in  the  other.  We  shall  then  have  the  equation 

3  Cu  +  8  HNO3  =  3  Cu(NO3)2  +  2  NO  +  4  H2O. 

A  number  of  other  reactions  may  take  place  when  nitric 
acid  acts  upon  metals,  resulting  in  the  formation  of  other 
oxides  of  nitrogen,  free  nitrogen,  or  even  ammonia.  The 
reaction  just  given  is,  however,  the  usual  one. 

Importance  of  steps  in  a  reaction.  This  complete  equation  has 
the  advantage  of  making  it  possible  to  calculate  very  easily  the  pro- 
portions in  which  the  various  substances  enter  into  the  reaction  or 
are  formed  in  it.  It  is  unsatisfactory  in  that  it  does  not  give  full 
information  about  the  way  in  which  the  reaction  takes  place.  For 
example,  it  does  not  suggest  that  hydrogen  is  at  first  formed,  and 
subsequently  transformed  into  water.  It  is  always  much  more  impor- 
tant to  remember  the  steps  in  a  chemical  reaction  than  to  remember 
the  equation  expressing  the  complete  action ;  for  if  these  steps  in 
the  reaction  are  understood,  the  complete  equation  is  easily  obtained 
in  the  manner  just  described. 

Salts  of  nitric  acid, — nitrates.  The  salts  of  nitric  acid 
are  called  nitrates.  Many  of  these  salts  will  be  described  in 
the  study  of  the  metals.  They  are  all  soluble  in  water, 
and  when  heated  to  a  high  temperature  undergo  decom- 
position. In  a  few  cases  a  nitrate  on  being  heated  evolves 
oxygen,  forming  a  nitrite : 

NaNO  =  NaNO  +  O. 


132     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

In  other  cases  the   decomposition  goes  further,  and  the 
metal  is  left  as  oxide  : 

Cu(NO3)2  =  CuO  +  2  NO2  +  O. 

Nitrous  acid  (HNO2).  It  is  an  easy  matter  to  obtain 
sodium  nitrite  (NaNO2),  as  the  reaction  given  on  the  previ- 
ous page  indicates.  Instead  of  merely  heating  the  nitrate,  it 
is  better  to  heat  it  together  with  a  mild  reducing  agent,  such 
as  lead,  when  the  reaction  takes  place  which  is  expressed  by 

the  equation    NaNQ3  +  Pb  =  PbO  +  NaNO2.    ''. 

When  sodium  nitrite  is  treated  with  an  acid,  such  as  sul- 
phuric acid,  it  is  decomposed  and  nitrous  acid  is  set  free : 

NaNO2  +  H2SO4  =  NaHSO4  +  HNO2. 

The  acid  is  very  unstable,  however,  and  decomposes  readily 
into  water  and  nitrogen  trioxide  (N2O3) : 

2  HNO2  =  H2O  +  N2O3. 
Dilute  solutions  of  the  acid,  however,  can  be  obtained. 

COMPOUNDS  OF  NITROGEN  WITH  OXYGEN 

Nitrogen  combines  with  oxygen  to  form  five  different 
oxides.  The  formulas  and  names  of  these  are  as  follows : 

N2O nitrous  oxide. 

NO nitric  oxide. 

NO2 nitrogen  peroxide. 

N2O3  .     .     .     4    .  nitrogen  trioxide,  or  nitrous  anhydride. 

N2O5 nitrogen  pentoxide,  or  nitric  anhydride. 

These  will  now  be  briefly  discussed. 

Nitrous  oxide  (laughing gas)  (N2O).  Ammonium  nitrate, 
like  all  nitrates,  undergoes  decomposition  when  heated ; 
and  owing  to  the  fact  that  it  contains  no  metal,  but  does 


COMPOUNDS  OF  NITROGEN 


133 


contain  both  oxygen  and  hydrogen,  the  reaction  is  a  peculiar 
one.    It  is  represented  by  the  equation 


The  oxide  of  nitrogen  so  formed  is  called  nitrous  oxide  or 
laughing  gas.  It  is  a  colorless  gas  having  a  slight  odor.  It 
is  somewhat  soluble  in  water,  and  in  solution  has  a  slightly 
sweetish  taste.  It  is  easily  converted  into  a  liquid  and  can 
be  purchased  in  this  form.  When  inhaled  it  produces  a  kind 
of  hysteria  (hence  the  name  "  laughing  gas  "),  and  even  un- 
consciousness and  insensibility  to  pain  if  taken 
in  large  amounts.  It  has  long  been  used  as  an 
anaesthetic  for  minor  surgical  operations,  such 
as  those  of  dentistry,  but  owing  to  its  un- 
pleasant after  effects  it  is  not  so  much  in  use 
now  as  formerly. 

Chemically,  nitrous  oxide  is  remarkable  for 
the  fact  that  it  is  a  very  energetic  oxidizing 
agent.  Substances  such  as  carbon,  sulphur, 
iron,  and  phosphorus  burn  in  it  almost  as 
brilliantly  as  in  oxygen,  forming  oxides  and 
setting  free  nitrogen.  Evidently  the  oxygen  in  nitrous 
oxide  cannot  be  held  in  verysfirm  combination  by  the 
nitrogen. 

Nitric  oxide  (NO).  We  have  seen  that  when  nitric  acid 
acts  upon  metals,  such  as  copper,  the  reaction  represented 
by  the  following  equation  takes  place  : 

3  Cu  +  8  HNO3  =  3  Cu(NO8)2  +  2  NO  -f-  4  H2O. 

Nitric  oxide  is  most  conveniently  prepared  in  this  way. 
The  metal  is  placed  in  the  flask  A  (Fig.  39)  and  the  acid 
added  slowly  through  the  funnel  tube  B.  The  gas  escapes 
through  C  and  is  collected  over  water. 


FIG.  39 


134     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Pure  nitric  oxide  is  a  colorless  gas,  slightly  heavier  than 
air,  and  is  practically  insoluble  in  water.  It  is  a  difficult 
gas  to  liquefy.  Unlike  nitrous  oxide,  nitric  oxide  does  not 
part  with  its  oxygen  easily,  and  burning  substances  intro- 
duced into  this  gas  are  usually  extinguished.  A  few  sub- 
stances like  phosphorus,  which  have  a  very  strong  affinity 
for  oxygen  and  which  are  burning  energetically  in  the  air, 
will  continue  to  burn  in  an  atmosphere  of  nitric  oxide.  In 
this  case  the  nitric  oxide  loses  all  of  its  oxygen  and  the 
nitrogen  is  set  free  as  gas. 

Action  of  nitric  oxide  with  oxygen.  When  nitric  oxide 
comes  into  contact  with  oxygen  or  with  the  air,  it  at  once 
combines  with  the  oxygen  even  at  ordinary  temperatures, 
forming  a  reddish-yellow  gas  of  the  formula  NO2,  which  is 
called  nitrogen  peroxide.  This  action  is  not  energetic  enough 
to  produce  a  flame,  though  considerable  heat  is  set  free. 

Nitrogen  peroxide  (NO2).  This  gas,  as  we  have  just  seen, 
is  formed  by  allowing  nitric  oxide  to  come  into  contact  with 
oxygen.  It  can  also  be  made  by  heating  certain  nitrates, 
such  as  lead  nitrate : 

Pb(N03)2  =  PbO  +  2  N02  +  O. 

It  is  a  reddish-yellow  gas  of  unpleasant  odor,  which  is  quite 
poisonous  when  inhaled.  It  is  heavier  than  air  and  is  easily 
condensed  to  a  liquid.  It  dissolves  in  water,  but  this  solu- 
tion is  not  a  mere  physical  solution;  the  nitrogen  peroxide 
is  decomposed,  forming  a  mixture  of  nitric  and  nitrous  acids : 

2  NO2  +  H2O  =  HNO2  +  HNO3. 

Nitrogen  peroxide  will  not  combine  with  more  oxygen ;  it 
will,  however,  give  up  a  part  of  its  oxygen  to  burning  sub- 
stances, acting  as  an  oxidizing  agent : 

NO9  =  NO  +  O. 


COMPOUNDS  OF  NITROGEN  135 

Acid  anhydrides.  The  oxides  N2O3  (nitrogen  trioxide) 
and  N2O5  (nitrogen  pentoxide)  are  rarely  prepared  and 
need  not  be  separately  described.  They  bear  a  very  inter- 
esting relation  to  the  acids  of  nitrogen.  When  dissolved  in 
water  they  combine  with  the  water,  forming  acids : 

N2O3-hH2O-2HNO2, 
N00R  +  H00  =  2  HNOc 


On  the  other  hand,  nitrous  acid  very  easily  decomposes, 
yielding  water  and  nitrogen  trioxide,  and  by  suitable  means 
nitric  acid  likewise,  may  be  decomposed  into  water  and 
nitrogen  pentoxide : 

2  HN02  =  H20  -h  N203, 
2  HN03  =  H20  +  N205. 

In  view  of  the  close  relation  between  these  oxides  and 
the  corresponding  acids,  they  are  called  anhydrides  of  the 
acids,  N2O3  being  nitrous  anhydride  and  N2O5  nitric 
anhydride. 

DEFINITION  :  Any  oxide  which  will  combine  with  water 
to  form  an  acid,  or  which  together  with  water  is  formed 
by  tJie  decomposition  of  an  acid,  is  called  an  anhydride  of 
that  acid. 

EXERCISES 

1.  Perfectly  dry  ammonia  does  not  affect  litmus  paper.    Explain. 

2.  Can  ammonia  be  dried  by  passing  the  gas  through  concen- 
trated sulphuric  acid  ?    Explain. 

3.  Ammonium    hydroxide   is    a  weak  base,  i.e.  it  is  not  highly 
dissociated.    When  it  is  neutralized  by  strong  acids  the  heat  of  reac- 
tion is  less  than  when  strong  bases  are  so  neutralized.    Suggest  some 
possible  cause  for  this. 

4.  Why  is  brine  used  in  the  manufacture  of  artificial  ice  ? 

5.  Discuss  the  energy  changes  which  take  place  in  the  manufac- 
ture of  artificial  ice. 


136     AN  ELEMENTARY  ST'UDY  OF  CHEMISTRY 

6.  What  weight  of  ammonium  chloride  is  necessary  to  furnish 
enough  ammonia  to  saturate  i  1.  of  water  at  o°  and  760  mm.? 

7.  What  weight  of  sodium  nitrate  is  necessary  to  prepare  100  cc. 
of  commercial  nitric   acid?    What  weight  of   potassium  nitrate  is 
necessary  to  furnish  the  same  weight  of  acid  ? 

8.  100  1.  of  nitrogen  peroxide  were  dissolved  in  water  and  neu- 
tralized with  sodium  hydroxide.    What  substances  were  formed  and 
how  much  of  each?     (i  1.  nitrogen  peroxide  weighs  2.05  grams.) 

9.  How  many  liters  of  nitrous  oxide,  measured  under  standard 
conditions,  can  be  prepared  from  10  g.  of  ammonium  nitrate? 

10.  What  weight  of  copper  is  necessary  to  prepare  50  1.  of  nitric 
oxide  under  standard  conditions  ? 

11.  (a)  Calculate  the  percentage   composition  of  the  oxides  of 
nitrogen.     (£)  What  important  law  does  this  series  of  substances 
illustrate  ? 

12.  Write  the  equations  representing  the  reactions  between  am- 
monium hydroxide,  and  sulphuric  acid  and  nitric  acid  respectively, 
in  accordance  with  the  theory  of  electrolytic  dissociation. 

13.  In  the  same  way,  write  the  equations  representing  the  reac- 
tions between  nitric  acid  and  each  of  the  following  bases  :  NaOH, 
KOH,  NH4OH,  Ca(OH)2. 


CHAPTER   XIII 
REVERSIBLE  REACTIONS  AND  CHEMICAL  EQUILIBRIUM 

Reversible  reactions.  The  reactions  so  far  considered 
have  been  represented  as  continuing,  when  once  started, 
until  one  or  the  other  substance  taking  part  in  the  reac- 
tion has  been  used  up.  In  some  reactions  this  is  not  the 
case.  For  example,  we  have  seen  that  when  steam  is  passed 
over  hot  iron  the  reaction  is  represented  by  the  equation 

3  Fe  +  4  H20  =  Fe304  +  8  H. 

On  the  other  hand,  when  hydrogen  is  passed  over  hot  iron 
oxide  the  reverse  reaction  takes  place  : 

Fe304  +  8  H  =  3  Fe  +  4  H2O. 

The  reaction  can  therefore  go  in  either  direction,  depending 
upon  the  conditions  of  the  experiment.  Such  a  reaction  is 
called  a  reversible  reaction.  It  is  represented  by  an  equa- 
tion with  double  arrows  in  place  of  the  equality  sign,  thus  : 

3  Fe  +  4  H20  ^  Fe3O4  +  8  H. 
In  a  similar  way,  the  equation  , 


expresses  the  fact  that  under  some  conditions  nitrogen  may 
unite  with  hydrogen  to  form  ammonia,  while  under  other 
conditions  ammonia  decomposes  into  nitrogen  and  hydrogen. 
The  conversion  of  oxygen  into  ozone  is  also  reversible 
and  may  be  represented  thus  : 

oxygen  ^  ozone. 


,  I 

138     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Chemical  equilibrium.  Reversible  reactions  do  not  usu- 
ally go  on  to  completion  in  one  direction  unless  the  condi- 
tions under  which  the  reaction  takes  place  are  very  carefully 
chosen.  Thus,  if  iron  and  steam  are  confined  in  a  heated 
tube,  the  steam  acts  upon  the  iron,  producing  iron  oxide 
and  hydrogen.  But  these  substances  in  turn  act  upon  each 
other  to  form  iron  and  steam  once  more.  When  these  two 
opposite  reactions  go  on  at  such  rates  that  the  weight  of 
the  iron  changed  into  iron  oxide  is  just  balanced  by  the 
weight  of  the  iron  oxide  changed  into  iron,  there  will  be 
no  further*  change  in  the  relative  weights  of  the  four  sub- 
stances present  in  the  tube.  The  reaction  is  then  said  to 
have  reached  an  equilibrium. 

Factors  which  determine  the  point  of  equilibrium.  There 
are  two  factors  which  have  a  great  deal  of  influence  in 
determining  the  point  at  which  a  given  reaction  will  reach 
equilibrium. 

1 .  Influence  of  the  chemical  nature  of  the  substances.    If 
two   reversible   reactions   of   the    same    general  kind   are 
selected,  it  has  been  found  that  the  point  of  equilibrium  is 
different  in  the  two  cases.    For  example,  in  the  reactions 
represented  by  the  equations 

3  Fe  +  4  H20  ^  Fe3O4  +  8  H, 
Zn  +  H2O  ^  ZnO  -f  2  H, 

the  equilibrium  will  be  reached  when  very  different  quan- 
tities of  the  iron  and  zinc  have  been  changed  into  oxides. 
The  individual  chemical  properties  of  the  iron  and  zinc 
have  therefore  marked  influence  upon  the  point  at  which 
equilibrium  will  be  reached. 

2.  Influence  of  relative  mass.    If  the  tube  in  which  the 
reaction  3  Fe  +  4  H2O  F^  Fe3O4  +  8  H 


REVERSIBLE  REACTIONS  139 

has  come  to  an  equilibrium  is  opened  and  more  steam  is 
admitted,  an  additional  quantity  of  the  iron  will  be  changed 
into  iron  oxide.  If  more  hydrogen  is  admitted,  some  of  the 
oxide  will  be  reduced  to  metal.  The  point  of  equilibrium 
is  therefore  dependent  upon  the  relative  masses  of  the 
substances  taking  part  in  the  reaction.  When  one  of  the 
substances  is  a  solid,  however,  its  mass  has  little  influence, 
since  it  is  only  the  extent  of  its  surface  which  can  affect 
the  reaction. 

Conditions  under  which  reversible  reactions  are  complete. 
If,  when  the  equilibrium  between  iron  and  steam  has  been 
reached,  the  tube  is  opened  and  a  current  of  steam  is 
passed  in,  the  hydrogen  is  swept  away  as  fast  as  it  is 
formed.  The  opposing  reaction  of  hydrogen  upon  iron 
oxide  must  therefore  cease,  and  the  action  of  steam  on 
the  iron  will  go  on  until  all  of  the  iron  has  been  trans- 
formed into  iron  oxide. 

On  the  other  hand,  if  a  current  of  hydrogen  is  admitted 
into  the  tube,  the  steam  will  be  swept  away  by  the  hydro- 
gen, and  all  of  the  iron  oxide  will  be  reduced  to  iron.  A 
reversible  reaction  can  therefore  be  completed  in  either  direc- 
tion when  one  of  the  products  of  the  reaction  is  removed  as 
fast  as  it  is  formed. 

Equilibrium  in  solution.  When  reactions  take  place  in 
solution  in  water  the  same  general  principles  hold  good. 
The  matter  is  not  so  simple,  however,  as  in  the  case  just 
described,  owing  to  the  fact  that  many  of  the  reactions  in 
solution  are  due  to  the  presence  of  ions.  The  substances 
most  commonly  employed  in  solution  are  acids,  bases,  or 
salts,  and  all  of  these  undergo  dissociation.  Any  equilib- 
rium which  may  be  reached  in  solutions  of  these  substances 
must  take  place  between  the  various  ions  formed,  on  the 


140     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

one  hand,  and  the  undissociated  molecules,  on  the  other. 
Thus,  when  nitric  acid  is  dissolved  in  water,  equilibrium  is 
reached  in  accordance  with  the  equation 

H+ +  NO^  HNOg. 

Conditions  under  which  reversible  reactions  in  solution 
are  complete.  The  equilibrium  between  substances  in 
solution  may  be  disturbed  and  the  reaction  caused  to  go 
on  in  one  direction  to  completion  in  either  of  three  ways. 

I.  A  gas  may  be  formed  which  escapes  from  the  solution. 
When  sodium  nitrate  and  sulphuric  acid  are  brought  to- 
gether in  solution  all  four  ions,  Na+,  NOg~,  H+,  SO^",  are 
formed.  These  ions  are  free  to  rearrange  themselves  in 
various  combinations.  For  example,  the  H+  and  the  NOg~ 
ions  will  reach  the  equilibrium 


If  the  experiment  is  performed  with  very  little  water 
present,  as  is  the  case  in  the  preparation  of  nitric  acid,  the 
equilibrium  will  be  reached  when  most  of  the  H+  and  the 
NOg~  ions  have  combined  to  form  undissociated  HNO3. 

Finally,  if  the  mixture  is  now  heated  above  the  boiling 
point  of  nitric  acid,  the  acid  distills  away  as  fast  as  it  is 
formed.  More  and  more  H+  and  NOg~  ions  will  then  com- 
bine, and  the  process  will  continue  until  one  or  the  other 
of  them  has  all  been  removed  from  the  solution.  The  sub- 
stance remaining  is  sodium  acid  sulphate  (NaHSO4),  and 
the  reaction  can  therefore  be  expressed  by  the  equation 

NaNOg  +  H2S04  =  NaHSO4  +  HNO3. 

2.  An  insoluble  solid  may  be  formed.  When  hydrochloric 
acid  (HC1)  and  silver  nitrate  (AgNO3)  are  brought  together 


REVERSIBLE  REACTIONS  141 

in  solution  the  following  ions  will  be  present  :  H+,  Cl~, 
Ag+,  NO^.    The  ions  Ag+  and  Cl~  will  then  set  up  the 

equilibrium 

- 


But  silver  chloride  (AgCl)  is  almost  completely  insoluble 
in  water,  and  as  soon  as  a  very  little  of  it  has  formed  the 
solution  becomes  supersaturated,  and  the  excess  of  the  salt 
precipitates.  More  silver  and  chlorine  ions  then  unite,  and 
this  continues  until  practically  all  of  the  silver  or  the  chlo- 
rine ions  have  been  removed  from  the  solution.  We  then 
say  that  the  following  reaction  is  complete  : 

AgNO8  +  HC1  =  AgCl  +  HNO3. 

3.  Two  different  ions  may  form  undissociated  molecules. 
In  the  neutralization  of  sodium  hydroxide  by  hydrochloric 
acid  the  ions  H+  and  OH~  come  to  the  equilibrium 


But  since  water  is  almost  entirely  undissociated,  equilib- 
rium can  only  be  reached  when  there  are  very  few  hydroxyl 
or  hydrogen  ions  present.  Consequently  the  two  ions  keep 
uniting  until  one  or  the  other  of  them  is  practically  removed 
from  the  solution.  When  this  occurs  the  neutralization 
expressed  in  the  following  equation  is  complete  : 

NaOH  +  HC1  =  H2O  +  NaCl. 

Preparation  of  acids.  The  principle  of  reversible,  reac- 
tions finds  practical  application  in  the  preparation  of  most 
of  the  common  acids.  An  acid  is  usually  prepared  by 
treating  the  most  common  of  its  salts  with  some  other 
acid  of  high  boiling  point.  The  mixture  is  then  heated 
until  the  lower  boiling  acid  desired  distills  out.  Owing  to 


142     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

its  high  boiling  point  (338°);  sulphuric  acid  is  usually 
employed  for  this  purpose,  most  other  acids  boiling  below 
that  temperature. 

EXERCISES 

1.  What  would  take  place  when  solutions  of  silver  nitrate  and 
sodium  chloride  are  brought  together  ?    What  other  chlorides  would 
act  in  the  same  way  ? 

2.  Is   the    reaction   expressed  by  the   equation    NH3  +  H2O  = 
NH4OH  reversible?    If  so,  state  the  conditions  under  which  it  will 
go  in  each  direction. 

3.  Is   the   reaction   expressed  by  the  equation  2  H  +  O  =  H2O 
reversible?    If  so,  state  the  conditions  under  which  it  will  go  in  each 
direction. 

4.  Suggest  a  method  for  the  preparation  of  hydrochloric  acid. 


CHAPTER  XIV 
SULPHUR  AND  ITS  COMPOUNDS 

Occurrence.  The  element  sulphur  has  been  known  from 
the  earliest  times,  since  it  is  widely  distributed  in  nature 
and  occurs  in  large  quantities  in  the  uncombined  form, 
especially  in  the  neighborhood  of  volcanoes.  Sicily  has 
long  been  famous  for  its  sulphur  mines,  and  smaller 
deposits  are  found  in  Italy,  Iceland,  Mexico,  and  espe- 
cially in  Louisiana,  where  it  is  mined  extensively.  In  com- 
bination, sulphur  occurs  abundantly  in  the  form  of  sulphides 
and  sulphates.  In  smaller  amounts  it  is  found  in  a  great 
variety  of  minerals,  and  it  is  a  constituent  of  many  animal 
and  vegetable  substances. 

Extraction  of  sulphur.  Sulphur  is  prepared  from  the 
native  substance,  the  separation  of  crude  sulphur  from 
the  rock  and  earthy  materials  with  which  it  is  mixed  being 
a  very  simple  process.  The  ore  from  the  mines  is  merely 
heated  until  the  sulphur  melts  and  drains  away  from  the 
earthy  impurities.  The  crude  sulphur  obtained  in  this  way 
is  distilled  in  a  retort-shaped  vessel  made  of  iron,  the  exit 
tube  of  which  opens  into  a  cooling  chamber  of  brickwork. 
When  the  sulphur  vapor  first  enters  the  cooling  chamber 
it  condenses  as  a  fine  crystalline  powder  called  flowers  of 
sulphur.  As  the  condensing  chamber  becomes  warm,  the 
sulphur  collects  as  a  liquid  in  it,  and  is  drawn  off  into 
cylindrical  molds,  the  product  being  called  roll  sulphur  or 
brimstone. 

M3 


144     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Physical  properties.  Roll  sulphur  is  a  pale  yellow,  crys- 
talline solid,  without  marked  taste  and  with  but  a  faint  odor. 
It  is  insoluble  in  water,  but  is  freely  soluble  in  a  few  liquids, 
notably  in  carbon  disulphide.  Roll  sulphur  melts  at  1 14.8°. 
Just  above  the  melting  point  it  forms  a  rather  thin,  straw- 
colored  liquid.  As  the  temperature  is  raised,  this  liquid 
turns  darker  in  color  and  becomes  thicker,  until  at  about 
235°  it  is  almost  black  and  is  so  thick  that  the  vessel  con- 
taining it  can  be  inverted  without  danger  of  the  liquid  run- 
ning out.  At  higher  temperatures  it  becomes  thin  once 
more,  and  boils  at  448°,  forming  a  yellowish  vapor.  On 
cooling  the  same  changes  take  place  in  reverse  order. 

Varieties  of  sulphur.  Sulphur  is  known  in  two  general 
forms,  crystalline  and  amorphous.  Each  of  these  forms 
exists  in  definite  modifications. 

Crystalline  sulphur.  Sulphur  occurs  in  two  crystalline 
forms,  namely,  rhombic  sulphur  and  monoclinic  sulphur. 

1 .  Rhombic  sulphur.    When  sulphur  crystallizes  from  its 
solution  in  carbon  disulphide  it  separates  in  crystals  which 
have  the  same  color  and  melting  point  as  roll  sulphur,  and 
are  rhombic  in  shape.    Roll  sulphur  is  made  up  of  minute 
rhombic  crystals. 

2.  Monoclinic  sulphur.    When  melted  sulphur  is  allowed 
to  cool  until  a  part   of  the  liquid  has  solidified,  and  the 
remaining  liquid  is   then  poured  off,  it  is  found  that  the 
solid  sulphur  remaining  in  the  vessel  has  assumed  the  form 
of  fine  needle-shaped  crystals.    These  differ  much  in  appear- 
ance from  the  rhombic  crystals  obtained   by  crystallizing 
sulphur  from  its  solution  in  carbon  disulphide.    The  needle- 
shaped  form  is  called  monoclinic  sulphur.    The  two  varieties 
differ  also  in  density  and  in  melting  point,  the  monoclinic 
sulphur  melting  at  120°. 


SULPHUR  AND  ITS  COMPOUNDS 


Monoclinic  and  rhombic  sulphur  remain  unchanged  in 
contact  with  each  other  at  96°.  Above  this  temperature  the 
rhombic  changes  into  monoclinic ;  at  lower  temperatures 
the  monoclinic  changes  into  rhombic.  The  temperature  96° 
is  therefore  called  the  transition  point  of  sulphur.  Heat  is 
set  free  when  monoclinic  sulphur  changes  into  rhombic. 

Amorphous  sulphur.  Two  varieties  of  amorphous  sul- 
phur can  be  readily  obtained.  These  are  white  sulphur 
and  plastic  sulphur. 

1 .  Wtiite  sulphur.   Flowers  of  sulphur,  the  preparation  of 
which  has  been  described,  consists  of  a  mixture  of  rhombic 
crystals   and    amorphous   par- 
ticles.   When  treated  with  car- 
bon disulphide,  the  crystals 

dissolve,  leaving  the  amorphous 
particles  as  a  white  residue. 

2.  Plastic  sulphur.     When 
boiling  sulphur  is  poured  into 
cold  water  it  assumes  a  gummy, 
doughlike  form,  which  is  quite 
elastic.    This  can  be   seen  in 
a  very  striking  manner  by  dis- 
tilling sulphur   from  a  small, 
short-necked  retort,  such  as  is 
represented    in    Fig.    40,   and 

allowing  the  liquid  to  run  directly  into  water.  In  a  few 
days  it  becomes  quite  brittle  and  passes  over  into  ordinary 
rhombic  sulphur. 

Chemical  properties  of  sulphur.  When  sulphur  is  heated 
to  its  kindling  temperature  in  oxygen  or  in  the  air  it  burns 
with  a  pale  blue  flame,  forming  sulphur  dioxide  (SO2).  Small 
quantities  of  sulphur  trioxide  (SO3)  may  also  be  formed  in 


FIG.  40 


146     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

the  combustion  of  sulphur.  Most  metals  when  heated  with 
sulphur  combine  directly  with  it,  forming  metallic  sulphides. 
In  some  cases  the  action  is  so  energetic  that  the  mass 
becomes  incandescent,  as  has  been  seen  in  the  case  of  iron 
uniting  with  sulphur.  This  property  recalls  the  action  of 
oxygen  upon  metals,  and  in  general  the  metals  which  com- 
bine readily  with  oxygen  are  apt  to  combine  quite  readily 
with  sulphur. 

Uses  of  sulphur.  Large  quantities  of  sulphur  are  used 
as  a  germicide  in  vineyards,  also  in  the  manufacture  of  gun- 
powder, matches,  vulcanized  rubber,  and  sulphuric  acid. 


COMPOUNDS  OF  SULPHUR  WITH  HYDROGEN 

Hydrosulphuric  acid  (H2S).  This  substance  is  a  gas  hav- 
ing the  composition  expressed  by  the  formula  H2S  and  is 
commonly  called  hydrogen  sulphide.  It  is  found  in  the 
vapors  issuing  from  volcanoes,  and  in  solution  in  the  so-called 
sulphur  waters  of  many  springs.  It  is  formed  when  organic 
matter  containing  sulphur  undergoes  decay,  just  as  ammonia 
is  formed  under  similar  circumstances  from 
nitrogenous  matter. 

Preparation.  Hydrosulphuric  acid  is  pre- 
pared in  the  laboratory  by  treating  a  sulphide 
with  an  acid.  Iron  sulphide  (FeS)  is  usually 
employed : 

FeS  +  2  HC1  =  FeCl2  +  H2S. 

A  convenient  apparatus  is  shown  in  Fig.  4 1 . 
A  few  lumps  of  iron  sulphide  are  placed  in  the 
FIG.  41         bottle  A,  and  dilute  acid  is  added  in  small 
quantities  at  a  time  through  the  funnel  tube  B>  the  gas 
escaping  through  the  tube  C. 


SULPHUR  AND  ITS  COMPOUNDS  147 

Explanation  of  the  reaction.  Iron  sulphide  is  a  salt  of  hydrosul- 
phuric  acid,  and  this  reaction  is  therefore  similar  to  the  one  which 
takes  place  when  sulphuric  acid  acts  upon  a  nitrate.  In  both  cases 
a  salt  and  an  acid  are  brought  together,  and  there  is  a  tendency  for 
the  reaction  to  go  on  until  a  state  of  equilibrium  is  reached.  This 
equilibrium  is  constantly  disturbed  by  the  escape  of  the  gaseous 
acid  set  free,  so  that  the  reaction  goes  on  until  all  of  the  original 
salt  has  been  decomposed.  The  two  reactions  differ  in  that  the  first 
one  is  complete  at  ordinary  temperatures,  while  in  the  case  of  sul- 
phuric acid  acting  upon  sodium  nitrate,  the  reacting  substances  must 
be  heated  so  as  to  secure  a  temperature  at  which  nitric  acid  is  a  gas. 

Physical  properties.  Hydrosulphuric  acid  is  a  colorless 
gas,  having  a  weak,  disagreeable  taste  and  an  exceedingly 
offensive  odor.  It  is  rather  sparingly  soluble  in  water  at 
ordinary  temperatures,  about  three  volumes  dissolving  in 
one  of  water.  In  boiling  water  it  is  not  soluble  at  all.  In 
pure  form  it  acts  as  a  violent  poison,  and  even  when  diluted 
largely  with  air  produces  headache,  dizziness,  and  nausea. 
It  is  a  little  heavier  than  air,  having  a  density  of  1.18. 

Chemical  properties.  The  most  important  chemical  prop- 
erties of  hydrosulphuric  acid  are  the  following: 

1.  Acid  properties.    Hydrosulphuric  acid  is  a  weak  acid. 
In  solution  in  water  it  turns  blue  litmus  red  and  neutralizes 
bases,  forming  salts  called  sulphides. 

2.  Action  on  oxygen.    The  elements  composing  hydro- 
sulphuric  acid  have  each  a  strong  affinity  for  oxygen,  and 
are  not  held  together  very  firmly.    Consequently  the  gas 
burns  readily  in  oxygen  or  the  air,  according  to  the  equation 

H2S  +  3  O  =  H20  +  S02. 

When  there  is  not  enough  oxygen  for  both  the  sulphur  and 
the  hydrogen,  the  latter  element  combines  with  the  oxygen 
and  the  sulphur  is  set  free: 

H2S  +  O  =  H2O  +  S. 

! 


148     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

3.  Reducing  action.    Owing  to  the  ease  with  which  hydro- 
sulphuric  acid  decomposes  and  the  strong  affinity  of  both 
sulphur  and  hydrogen  for  oxygen,  the  substance  is  a  strong 
reducing  agent,  taking  oxygen  away  from  many  substances 
which  contain  it. 

4.  Action  on  metals.    Hydrosulphuric  acid  acts  towards 
metals  in  a  way  very  similar  to  water.    Thus,  when  it  is 
passed  over  heated  iron  in  a  tube,  the  reaction  is  represented 
by  the  equation 

3  Fe  +  4  H2S  =  Fe3  S4  +  8  H. 

Water  in  the  form  of  steam,  under  similar  circumstances, 
acts  according  to  the  equation 

Salts  of  hydrosulphuric  acid, —  sulphides.  The  salts  oj 
hydrosulphuric  acid,  called  sulphides,  form  an  important 
class  of  salts.  Many  of  them  are  found  abundantly  in 
nature,  and  some  of  them  are  important  ores.  They  will 
be  frequently  mentioned  in  connection  with  the  metals. 

Most  of  the  sulphides  are  insoluble  in  water,  and  some 
of  them  are  insoluble  in  acids.  Consequently,  when  hydro- 
sulphuric  acid  is  passed  into  a  solution  of  a  salt,  it  often 
happens  that  a  sulphide  is  precipitated.  With  copper  chloride 
the  equation  is 

CuCl2  +  H2S  =  CuS  +  2  HC1. 

Because  of  the  fact  that  some  metals  are  precipitated  in 
this  way  as  sulphides  while  others  are  not,  hydrosulphuric 
acid  is  extensively  used  in  the  separation  of  the  metals  in 
the  laboratory. 

Explanation  of  the  reaction.  When  hydrosulphuric  acid  and  copper 
chloride  are  brought  together  in  solution,  both  copper  and  sulphur 


SULPHUR  AND  ITS  COMPOUNDS  149 

ions  are  present,  and  these  will  come  to  an  equilibrium,  as  represented 

in  the  equation 

Cu+  +  S~^  CuS. 

Since  copper  sulphide  is  almost  insoluble  in  water,  as  soon  as  a  very 
small  quantity  has  formed  the  solution  becomes  supersaturated,  and 
the  excess  keeps  precipitating  until  nearly  all  the  copper  or  sulphur 
ions  have  been  removed  from  the  solution.  With  some  other  ions, 
such  as  iron,  the  sulphide  formed  does  not  saturate  the  solution,  and 
no  precipitate  results. 

OXIDES  OF  SULPHUR 

Sulphur  forms  two  well-known  compounds  with  oxygen  : 
sulphur  dioxide  (SO2),  sometimes  called  sulphurous  anhy- 
dride ;  and  sulphur  trioxide  (SO3),  frequently  called  sul- 
phuric anhydride. 

Sulphur  dioxide  (SO2).  Sulphur  dioxide  occurs  in  nature 
in  the  gases  issuing  from  volcanoes,  and  in  solution  in  the 
water  of  many  springs.  It  is  likely  to  be  found  wherever 
sulphur  compounds  are  undergoing  oxidation. 

Preparation.  Three  general  ways  may  be  mentioned  for 
the  preparation  of  sulphur  dioxide : 

1.  By  the   combustion    of  sulphur.    Sulphur  dioxide  is 
readily  formed  by  the  combustion  of  sulphur  in  oxygen  or 

the  air  :  S  +  2  O  =  S02. 

It  is  also  formed  when  substances  containing  sulphur  are 

burned  :  ZnS  +  3  O  =  ZnO  +  SO2. 

2.  By  the  reduction  of  sulphuric  acid.    When  concentrated 
sulphuric  acid  is  heated  with  certain  metals,  such  as  copper, 
part  of  the  acid  is  changed  into  copper  sulphate,  and  part 
is  reduced  to  sulphurous  acid.    The  latter  then  decomposes 
into  sulphur  dioxide  and  water,  the  complete  equation  being 

Cu  +  2  H2SO4  =  CuSO4  +  SO2  +  2  H2O. 


150     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

3.  By  the  action  of  an  acid  on  a  sulphite.  Sulphites  are 
salts  of  sulphurous  acid  (H2SO3) .  When  a  sulphite  is  treated 
with  an  acid,  sulphurous  acid  is  set  free,  and  being  very 
unstable,  decomposes  into  water  and  sulphur  dioxide.  These 
reactions  are  expressed  in  the  equations 

Na2SO3  +  2  HC1  =  2  NaCl  +  H2SO3, 
H2SO3  =  H2O  +  SO2. 

Explanation  of  the  reaction.  In  this  case  we  have  two  reversible 
reactions  depending  on  each  other.  In  the  first  reaction,  . 

(i)        Na2SO3  +  2HC1  ^  2NaCl  +  H2SO3, 

we  should  expect  an  equilibrium  to  result,  for  none  of  the  four 
substances  in  the  equation  are  insoluble  or  volatile  when  water  is 
present  to  hold  them  in  solution.  But  the  quantity  of  the  H2SO3 
is  constantly  diminishing,  owing  to  the  fact  that  it  decomposes,  as 
represented  in  the  equation 

(2)        H2S03^H2O  +  S02, 

and  the  sulphur  dioxide,  being  a  gas,  escapes.  No  equilibrium  can 
therefore  result,  since  the  quantity  of  the  sulphurous  acid  is  constantly 
being  diminished  because  of  the  escape  of  sul- 
phur dioxide. 

Physical  properties.  Sulphur  dioxide  is 
a  colorless  gas,  which  at  ordinary  temper- 
atures is  2.2  times  as  heavy  as  air.  It  has 
a  peculiar,  irritating  odor.  The  gas  is  very 
soluble  in  water,  one  volume  of  water  dis- 
solving eighty  of  the  gas  under  standard 
conditions.  It  is  easily  condensed  to  a 
colorless  liquid,  and  can  be  purchased  in 
this  condition  stored  in  strong  bottles, 
FlG-42  such  as  the  one  represented  in  Fig.  42. 

Chemical  properties.  Sulphur  dioxide  has  a  marked  tend- 
ency to  combine  with  other  substances,  and  is  therefore  an 


SULPHUR  AND  ITS  COMPOUNDS  151 

active  substance  chemically.  It  combines  with  oxygen  gas, 
but  not  very  easily.  It  can,  however,  take  oxygen  away  from 
some  other  substances,  and  is  therefore  a  good  reducing 
agent.  Its  most  marked  chemical  property  is  its  ability  to 
combine  with  water  to  form  sulphurous  acid  (H2SO3). 

Sulphurous  acid  (H2SO3).  When  sulphur  dioxide  dissolves 
in  water  it  combines  chemically  with  it  to  form  sulphurous 
acid,  an  unstable  substance  having  the  formula  H2SO3. 
It  is  impossible  to  prepare  this  acid  in  pure  form,  as  it 
breaks  down  very  easily  into  water  and  sulphur  dioxide. 
The  reaction  is  therefore  reversible,  and  is  expressed  by 
the  equation  J^Q  +  ^  _^  HZSO3. 

Solutions  of  the  acid  in  water  have  a  number  of  interesting 
properties. 

1.  Acid  properties.    The  solution  has  all  the  properties 
typical  of  an  acid.    When  neutralized  by  bases,  sulphurous 
acid  yields  a  series  of  salts  called  sulphites. 

2.  Reducing  properties.    Solutions  of  sulphurous  acid  act 
as  good  reducing  agents.    This  is  due  to  the  fact  that  sul- 
phurous acid  has  the  power  of  taking  up  oxygen  from  the 
air,  or  from  substances  rich  in  oxygen,  and  is  changed  by 
this  reaction  into  sulphuric  acid : 

H2SO3  +  O  =  H2SO4, 
H2S03  +  H202  =  H2S04  +  H20. 

3.  Bleaching   properties.    Sulphurous    acid    has    strong 
bleaching  properties,  acting  upon  many  colored  substances 
in  such  a  way  as  to  destroy  their  color.     It  is  on  this  account 
used  to  bleach  paper,  straw  goods,  and  even  such  foods  as 
canned  corn. 

4.  Antiseptic  properties.  -  Sulphurous  acid    has   marked 
antiseptic  properties,  and  on  this  account  has  the  power 


152     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


of  arresting  fermentation.  It  is  therefore  used  as  a  pre- 
servative. 

Salts  of  sulphurous  acid,  —  sulphites.  The  sulphites,  like 
sulphurous  acid,  have  the  power  of  taking  up  oxygen  very 
readily,  and  are  good  reducing  agents.  On  account  of  this 
tendency,  commercial  sulphites  are  often  contaminated  with 
sulphates.  A  great  deal  of  sodium  sulphite  is  used  in  the 
bleaching  industry,  and  as  a  reagent  for  softening  paper  pulp. 

Sulphur  trioxide  (SO3).  When  sulphur  dioxide  and  oxy- 
gen are  heated  together  at  a  rather  high  temperature,  a 
small  amount  of  sulphur  trioxide  (SO3)  is  formed,  but  the 
reaction  is  slow  and  incomplete.  If,  however,  the  heating 
takes  place  in  the  presence  of  very  fine  platinum  dust,  the 
reaction  is  rapid  and  nearly  complete. 

Experimental  preparation  of  sulphur  trioxide.  The  experiment  can 
be  performed  by  the  use  of  the  apparatus  shown  in  Fig.  43,  the  fine 
platinum  being  secured  by  moistening  asbestos  fiber  with  a  solution 
of  platinum  chloride  and  igniting  it  in  a  flame.  The  fiber,  covered 
with  fine  platinum,  is  placed  in  a  tube  of  hard  glass,  which  is  then 


FIG. 43 

heated  with  a  burner  to  about  350°,  while  sulphur  dioxide  and  air  are 
passed  into  the  tube.  Union  takes  place  at  once,  and  the  strongly 
fuming  sulphur  trioxide  escapes  from  the  jet  at  the  end  of  the  tube, 
and  may  be  condensed  by  surrounding  the  receiving  tube  with  a 
freezing  mixture. 

Properties  of  sulphur  trioxide.    Sulphur  trioxide  is  a  col- 
orless liquid,  which  solidifies  at  about  15°  and  boils  at  46°. 


SULPHUR  AND  ITS  COMPOUNDS  153 

A  trace  of  moisture  causes  it  to  solidify  into  a  mass  of 
silky  white  crystals,  somewhat  resembling  asbestos  fiber 
in  appearance.  In  contact  with  the  air  it  fumes  strongly, 
and  when  thrown  upon  water  it  dissolves  with  a  hissing  'sound 
and  the  liberation  of  a  great  deal  of  heat.  The  product  of 
this  reaction  is  sulphuric  acid,  so  that  sulphur  trioxide  is 
the  anhydride  of  that  acid  : 

S03+H20  =  H2S04. 

Catalysis.  It  has  been  found  that  many  chemical  reac- 
tions, such  as  the  union  of  sulphur  dioxide  with  oxygen,  are 
much  influenced  by  the  presence  of  substances  which  do 
not  themselves  seem  to  take  a  part  in  the  reaction,  and  are 
left  apparently  unchanged  after  it  has  ceased.  These  reac- 
tions go  on  very  slowly  under  ordinary  circumstances,  but  are 
greatly  hastened  by  the  presence  of  the  foreign  substance. 
Substances  which  hasten  very  slow  reactions  in  this  way 
are  said  to  act  as  catalytic  agents  or  catalyzers,  and  the 
action  is  called  catalysis.  Just  how  the  action  is  brought 
about  is  not  well  understood. 

DEFINITION  :  A  catalyzer  is  a  substance  which  changes 
the  velocity  of  a  reaction^  but  does  not  change  its  products. 

Examples  of  Catalysis.  We  have  already  had  several 
instances  of  such  action.  Oxygen  and  hydrogen  combine 
with  each  other  at  ordinary  temperatures  in  the  presence 
of  platinum  powder,  while  if  no  catalytic  agent  is  present 
they  do  not  combine  in  appreciable  quantities  until  a  rather 
high  temperature  is  reached.  Potassium  chlorate,  when 
heated  with  manganese  dioxide,  gives  up  its  oxygen  at  a 
much  lower  temperature  than  when  heated  alone.  Hydrogen 
dioxide  decomposes  very  rapidly  when  powdered  manganese 
dioxide  is  sifted  into  its  concentrated  solution. 


154 


AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


On  the  other  hand,  the  catalytic  agent  sometimes  retards 
chemical  action .  For  example,  a  solution  of  hydrogen  dioxide 
decomposes  more  slowly  when  it  contains  a  little  phosphoric 
acid  than  when  perfectly  pure.  For  this  reason  commercial 
hydrogen  dioxide  always  contains  phosphoric  acid. 

Many  reactions  are  brought  about  by  the  catalytic  action 
of  traces  of  water.  For  example,  phosphorus  will  not  burn 
in  oxygen  in  the  absence  of  all  moisture.  Hydrochloric 
acid  will  not  unite  with  ammonia  if  the  reagents  are  per- 
fectly dry.  It  is  probable  that  many  of  the  chemical  trans- 
formations in  physiological  processes,  such  as  digestion,  are 
assisted  by  certain  substances  acting  as  catalytic  agents. 
The  principle  of  catalysis  is  therefore  very  important. 

Sulphuric  acid  (oil  of  vitriol)  (H2SO4).  Sulphuric  acid  is 
one  of  the  most  important  of  all  manufactured  chemicals. 
Not  only  is  it  one  of  the  most  common  reagents  in  the 
laboratory,  but  enormous  quantities  of  it  are  used  in  many 
of  the  industries,  especially  in  the  refining  of  petroleum, 
the  manufacture  of  nitroglycerin,  sodium  carbonate,  and 
fertilizers. 

Manufacture  of  sulphuric  acid.  i.  Contact  process.  The 
reactions  taking  place  in  this  process  are  represented  by 
the  following  equations  : 

.  SO2  +  O  =  SO3, 
S03+H20  =  H2S04. 

To  bring  about  the  first  of  these  reactions  rapidly,  a  cata- 
lyzer is  employed,  and  the  process  is  carried  out  in  the 
following  way:  Large  iron  tubes  are  packed  with  some 
porous  material,  such  as  calcium  and  magnesium  sulphates, 
which  contains  a  suitable  catalytic  substance  scattered 
through  it.  The  catalyzers  most  used  are  platinum  powder, 


SULPHUR  AND  ITS  COMPOUNDS  155 

vanadium  oxide,  and  iron  oxide.  Purified  sulphur  dioxide 
and  air  are  passed  through  the  tubes,  which  are  kept  at  a 
temperature  of  about  350°.  Sulphur  trioxide  is  formed,  and 
as  it  issues  from  the  tube  it  is  absorbed  in  water  or  dilute 
sulphuric  acid.  The  process  is  continued  until  all  the  water 
in  the  absorbing  vessel  has  been  changed  into  sulphuric  acid, 
so  that  a  very  concentrated  acid  is  made  in  this  way.  An 
excess  of  the  trioxide  may  dissolve  in  the  strong  sulphuric 
acid,  forming  what  is  known  as  fuming  stdphuric  acid. 

2.  Chamber  process.  The  method  of  manufacture  exclu- 
sively employed  until  recent  years,  and  still  in  very  exten- 
sive use,  is  much  more  complicated.  The  reactions  are 
quite  involved,  but  the  conversion  of  water,  sulphur  dioxide, 
and  oxygen  into  sulphuric  acid  is  accomplished  by  the  cata- 
lytic action  of  oxides  of  nitrogen.  The  reactions  are  brought 
about  in  large  lead-lined  chambers,  into  which  oxides  of 
nitrogen,  sulphur  dioxide,  steam,  and  air  are  introduced  in 
suitable  proportions. 

Reactions  of  the  chamber  process.  In  a  very  general  way,  the 
various  reactions  which  take  place  in  the  lead  chambers  may  be 
expressed  in  two  equations.  In  the  first  reaction  sulphur  dioxide, 
nitrogen  peroxide,  steam,  and  oxygen  unite,  as  shown  in  the  equation 

(1)  2  SO2  +  2  N02  +  H2O  +  O  =  2  SO2  (OH)  (NO2). 

The  product  formed  in  this  reaction  is  called  nitrosulphuric  acid 
or  "  chamber  crystals."  It  actually  separates  on  the  walls  of  the  cham- 
bers when  the  process  is  not  working  properly.  Under  normal  con- 
ditions, it  is  decomposed  as  fast  as  it  is  formed  by  the  action  of 
excess  of  steam,  as  shown  in  the  equation 

(2)  2  SO2  (OH)  (NO2)  +  H2O  +  O  =  2  H2S04  +  2  NO2. 

The  nitrogen  dioxide  formed  in  this  reaction  can  now  enter  into 
combination  with  a  new  quantity  of  sulphur  dioxide,  steam,  and  oxy- 
gen, and  the  series  of  reactions  go  on  indefinitely.  Many  other 
reactions  occur,  but  these  two  illustrate  the  principle  of  the  process. 


156     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


The  relation  between  sulphuric  acid  and  nitrosulphuric  acid  can 
be  seen  by  comparing  their  structural  formulas : 


/OH 


/OH 


The  latter  may  be  regarded  as  derived  from  the  former  by  the  sub- 
stitution of  the  nitro  group  (NO2)  for  the  hydroxyl  group  (OH). 

The  sulphuric  acid  plant,  Fig.  44  illustrates  the  simpler  parts  of 
a  plant  used  in  the  manufacture  of  sulphuric  acid  by  the  chamber 
process.  Sulphur  or  some  sulphide,  as  FeS2,  is  burned  in  furnace  A. 
The  resulting  sulphur  dioxide,  together  with  air  and  some  nitrogen 
peroxide,  are  conducted  into  the  large  chambers,  the  capacity  of  each 
^r  chamber  being  about  75,000  cu.  ft.  Steam  is  also  ad- 

mitted  into  these  chambers  at  different  points.    These 
compounds  react  to  form  sulphuric  acid,  according  to 


Glover 
tower 


B  A 


the  equations  given  above.  The  nitrogen  left  after  the  withdrawal  of 
the  oxygen  from  the  admitted  air  escapes  through  the  Gay-Lussac 
tower  X.  In  order  to  prevent  the  escape  of  the  oxides  of  nitrogen 
regenerated  in  the  reaction,  the  tower  is  filled  with  lumps  of  coke, 
over  which  trickles  concentrated  sulphuric  acid  admitted  from  Y. 
The  nitrogen  peroxide  dissolves  in  the  acid  and  the  resulting  solution 
collects  in  H.  This  is  pumped  into  £",  where  it  is  mixed  with  dilute 
acid  and  allowed  to  trickle  down  through  the  chamber  D  (Glover 
tower),  which  is  filled  with  some  acid-resisting  rock.  Here  the  nitro- 
gen peroxide  is  expelled  from  the  solution  by  the  action  of  the  hot 
gases  entering  from  A,  and  together  with  them  enters  the  first  chamber 
again.  The  acid  from  which  the  nitrogen  peroxide  is  expelled  col- 
lects in  F.  Theoretically,  a  small  amount  of  nitrogen  peroxide  would 
suffice  to  prepare  an  unlimited  amount  of  sulphuric  acid  ;  practically, 
some  of  it  escapes,  and  this  is  replaced  by  small  amounts  admitted  at  B. 


Lead-lined  chambers 
FIG.  44 


SULPHUR  AND  ITS  COMPOUNDS  157 

The  sulphuric  acid  so  formed,  together  with  the  excess  of  con- 
densed steam,  collect  upon  the  floor  of  the  chambers  in  the  form 
of  a  liquid  containing  from  62%  to  70%  of  sulphuric  acid.  The  product 
is  called  chamber  acid  and  is  quite  impure ;  but  for  many  purposes, 
such  as  the  manufacture  of  fertilizers,  it  needs  no  further  treatment. 
It  can  be  concentrated  by  boiling  it  in  vessels  made  of  iron  or  plati- 
num, which  resist  the  action  of  the  acid,  nearly  all  the  water  boiling 
off.  Pure  concentrated  acid  can  be  made  best  by  the  contact  process, 
while  the  chamber  process  is  cheaper  for  the  dilute  impure  acid. 

Physical  properties.  Sulphuric  acid  is  a  colorless,  oily 
liquid,  nearly  twice  as  heavy  as  water.  The  ordinary  con- 
centrated acid  contains  about  2  %  of  water,  has  a  density  of 
i  .84,  and  boils  at  338°.  It  is  sometimes  called  oil  of  -vitriol, 
since  it  was  formerly  made  by  distilling  a  substance  called 
green  vittiol. 

Chemical  properties.  Sulphuric  acid  possesses  chemical 
properties  which  make  it  one  of  the  most  important  of 
chemical  substances. 

1 .  Action  as  an  acid.    In  dilute  solution  sulphuric  acid 
acts   as    any  other  acid,    forming   salts  with    oxides   and 
hydroxides. 

2.  Action  as  an  oxidizing  agent.    Sulphuric  acid  contains 
a  large  percentage  of  oxygen  and  is,  like  nitric  acid,  a  very 
good    oxidizing    agent.    When    the    concentrated    acid    is 
heated  with  sulphur,  carbon,  and  many^  other  substances, 
oxidation    takes    place,    the    sulphuric    acid    decomposing 
according  to  the  equation 

H2SO4  =  H2SO3+O. 

3.  Action  on  metals :  In  dilute  solution  sulphuric  acid  acts 
upon  many  metals,  such  as  zinc,  forming  a  sulphate  and 
liberating  hydrogen.    When  the  concentrated  acid  is  em- 
ployed the  hydrogen  set  free  is  oxidized  by  a  new  portion 


158     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

of  the  acid,  with  the  liberation  of  sulphur  dioxide.    With 
copper  the  reactions  are  expressed  by  the  equations 

(1)  Cu  4-  H2SO4  =  CuSO4  +  2  H, 

(2)  H2S04  +  2  H  =  H2S03  +  H20, 

(3)  H2SO3  =  H2O  +  SO2. 

By  combining  these  equations  the  following  one  is  obtained  : 
Cu  4-  2  H2SO4  =  CuSO4  +  SO2  +  2  H2O. 

4.  Action  on  salts.    We  have  repeatedly  seen  that   an 
acid  of  high  boiling  point  heated  with  the  salt  of  some  acid 
of  lower  boiling  point  will  drive  out  the  low  boiling  acid. 
The  boiling  point  of  sulphuric  acid  (338°)  is  higher  than 
that  of  almost  any  common  acid ;  hence  it  is  used  largely 
in  the  preparation  of  other  acids. 

5.  Action  on  water.    Concentrated  sulphuric  acid  has  a 
very  great  affinity  for  water,  and  is  therefore  an  effective 
dehydrating  agent.    Gases  which  have  no  chemical  action 
upon  sulphuric  acid  can  be  freed  from  water  vapor  by  bub- 
bling them  through  the   strong  acid.    When  the  acid  is 
diluted  with  water  much  heat  is  set  free,  and  care  must 
be  taken  to  keep  the  liquid  thoroughly  stirred  during  the 
mixing,  and  to  pour  the  acid  into  the  water,  —  never  the 
reverse. 

Not  only  can  sulphuric  acid  absorb  water,  but  it  will 
often  withdraw  the  elements  hydrogen  and  oxygen  from  a 
compound  containing  them,  decomposing  the  compound, 
and  combining  with  the  water  so  formed.  For  this  reason 
most  organic  substances,  such  as  sugar,  wood,  cotton,  and 
woolen  fiber,  and  even  flesh,  all  of  which  contain  much 
oxygen  and  hydrogen  in  addition  to  carbon,  are  charred  or 
burned  by  the  action  of  the  concentrated  acid. 


SULPHUR  AND  ITS  COMPOUNDS  159 

Salts  of  sulphuric  acid, —  sulphates.  The  sulphates  form 
a  very  important  class  of  salts,  and  many  of  them  have  com- 
mercial uses.  Copperas  (iron  sulphate),  blue  vitriol  (copper 
sulphate),  and  Epsom  salt  (magnesium  sulphate)  serve  as 
examples.  Many  sulphates  are  important  minerals,  promi- 
nent among  these  being  gypsum  (calcium  sulphate)  and 
barytes  (barium  sulphate). 

Thiosulphuric  acid  (H2S2O3) ;  Thiosulphates.  Many  other  acids 
of  sulphur  containing  oxygen  are  known,  but  none  of  them  are  of 
great  importance.  Most  of  them  cannot  be  prepared  in  a  pure  state, 
and  are  known  only  through  their  salts.  The  most  important  of 
these  is  thiosulphuric  acid. 

When  sodium  sulphite  is  boiled  with  sulphur  the  two  substances 
combine,  forming  a  salt  which  has  the  composition  represented  in  the 
formula  Na2S2O3  :  ^^  +  g  =  ^5^. 

The  substance  is  called  sodium  thiosulphate,  and  is  a  salt  of  the 
easily  decomposed  acid  H2S2O3,  called  thiosulphuric  acid.  This 
reaction  is  quite  similar  to  the  action  of  oxygen  upon  sulphites : 

Na2SO3  +  O  =  Na2SO4. 

More  commonly  the  salt  is  called  sodium  hyposulphite,  or  merely 
"hypo."  It  is  a  white  solid  and  is  extensively  used  in  photography, 
in  the  bleaching  industry,  and  as  a  disinfectant. 

Monobasic  and  dibasic  acids.  Such  acids  as  hydrochloric 
and  nitric  acids,  which  have  only  one  replaceable  hydrogen 
atom  in  the  molecule,  or  in  other  words  yield  one  hydrogen 
ion  in  solution,  are  called  monobasic  acids.  Acids  yielding 
two  hydrogen  ions  in  solution  are  called  dibasic  acids. 
Similarly,  we  may  have  tribasic  and  tetrabasic  acids.  The 
three  acids  of  sulphur  are  dibasic  acids.  It  is  therefore  pos- 
sible for  each  of  them  to  form  both  normal  and  acid  salts. 
The  acid  salts  can  be  made  in  two  ways :  the  acid  may  be 
treated  with  only  half  enough  base  to  neutralize  it,  — 

NaOli  +  H2SO4  =  NariSQ4  +  H2O  ; 


160     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


or  a  normal  salt  may  be  treated  with  the  free  acid,  — 
Na2SO4  +  H2SO4  -  2  NaHSO4. 

Acid  sulphites  and   sulphides  may  be  made  in  the  same 
ways. 

Carbon  disulphide  (CS2).  When  sulphur  vapor  is  passed 
over  highly  heated  carbon  the  two  elements  combine,  form- 
ing carbon  disulphide  (CS2),  just  as  oxygen  and  carbon  unite 
to  form  carbon  dioxide  (CO2).  The  substance  is  a  heavy, 
colorless  liquid,  possessing,  when  pure,  a  pleasant  ethereal 
odor.  On  standing  for  some  time,  especially  when  exposed 
to  sunlight,  it  undergoes  a  slight  decomposition  and  acquires 
a  most  disagreeable,  rancid  odor.  It  has  the  property  of 
dissolving  many  substances,  such  as  gums,  resins,  and  waxes, 
which  are  insoluble  in  most  liquids,  and  it  is  extensively 
used  as  a  solvent  for  such  substances.  It  is  also  used  as  an 
insecticide.  It  boils  at  a  low  temperature  (46°),  and  its  vapor 

is  very  inflammable,  burn- 
ing in  the  air  to  form  car- 
bon dioxide  and  sulphur 
dioxide,  according  to  the 
equation 

2SO. 


FIG.  45 


Commercial  preparation  of 
carbon  disulphide.  In  the  prep- 
aration of  carbon  disulphide 
an  electrical  furnace  is  em- 
ployed, such  as  is  represented  in 
Fig.  45.  The  furnace  is  packed 
with  carbon  C,  and  this  is  fed 


in  through  the  hoppers  B,  as  fast  as  that  which  is  present  in  the 
hearth  of  the  furnace  is  used  up.  Sulphur  is  introduced  at  A,  and  at 
the  lower  ends  of  the  tubes  it  is  melted  by  the  heat  of  the  furnace  and 


SULPHUR  AND  ITS  COMPOUNDS  161 

flows  into  the  hearth  as  a  liquid.  An  electrical  current  is  passed 
through  the  carbon  and  melted  sulphur  from  the  electrodes  E,  heat- 
ing the  charge.  The  vapors  of  carbon  disulphide  pass  up  through 
the  furnace  and  escape  at  D,  from  which  they  pass  to  a  suitable 
condensing  apparatus. 

Comparison  of  sulphur  and  oxygen.  A  comparison  of 
the  formulas  and  the  chemical  properties  of  corresponding 
compounds  of  oxygen  and  sulphur  brings^^AMffkf  many 
striking  similarities.  The  conduct  of  hydrRft^munc  acid 
and  water  toward  many  substances  has  been  seen  to  be  very 
similar  ;  the  oxides  and  sulphides  of  the  metals  have  analo- 
gous formulas  and  undergo  many  parallel  reactions.  Carbon 
dioxide  and  disulphide  are  prepared  in  similar  ways  and 
undergo  many  analogous  reactions.  It  is  clear,  therefore, 
that  these  two  elements  are  far  more  closely  related  to  each 
other  than  to  any  of  the  other  elements  so  far  studied. 

Selenium  and  tellurium.  These  two  very  uncommon  elements  are 
still  more  closely  related  to  sulphur  than  is  oxygen.  They  occur  in 
comparatively  small  quantities  and  are  usually  found  associated  with 
sulphur  and  sulphides,  either  as  the  free  elements  or  more  commonly 
in  combination  with  metals.  They  form  compounds  with  hydrogen 
of  the  formulas  H2Se  and  H2Te ;  these  bodies  are  gases  with  prop- 
erties very  similar  to  those  of  H2S.  They  also  form  oxides  and 
oxygen  acids  which  resemble  the  corresponding  sulphur  compounds. 
The  elements  even  have  allotropic  forms  corresponding  very  closely 
to  those  of  sulphur.  Tellurium  is  sometimes  found  in  combination 
with  gold  and  copper,  and  occasions  some  difficulties  in  the  refining 
of  these  metals.  The  elements  have  very  few  practical  applications. 

Crystallography.  In  order  to  understand  the  difference  between 
the  two  kinds  of  sulphur  crystals,  it  is  necessary  to  know  something 
about  crystals  in  general  and  the  forms  which  they  may  assume. 
An  examination  of  a  large  number  of  crystals  has  shown  that  although 
they  may  differ  much  in  geometric  form,  they  can  all  be  considered 
as  modifications  of  a  few  simple  plans.  The  best  way  to  understand 
the  relation  of  one  crystal  to  another  is  to  look  upon  every  crystal 
as  having  its  faces  and  angles  arranged  in  definite  fashion  about 


1 62     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


certain  imaginary  lines  drawn  through  the  crystal.  These  lines  are 
called  axes,  and  bear  much  the  same  relation  to  a  crystal  as  do  the 
axis  and  parallels  of  latitude  and  longitude  to  the  earth  and  a  geo- 
graphical study  of  it.  All  crystals  can  be  referred  to  one  of  six 
simple  plans  or  systems,  which  have  their  axes  as  shown  in  the 
following  drawings. 

The  names  and  characteristics  of  these  systems  are  as  follows  : 

1.  Isometric  or  regular  system  (Fig.  46).    Three  equal  axes,  all 
at  right  angles. 

2.  Tetragonal  system  (Fig.  47).    Two  equal  axes  and  one  of  dif- 
ferent length,  all  at  right  angles  to  each  other. 


FIG.  46 


FIG.  47 


FIG.  48 


FIG.  49 


FIG.  50 


FIG.  51 


3.  Orthorhombic  system  (Fig.  48).    Three  unequal  axes,  all  at 
right  angles  to  each  other. 

4.  Monoclinic  system  (Fig.  49).    Two  axes  at  right  angles,  and 
a  third  at  right  angles  to  one  of  these,  but  inclined  to  the  other. 

5.  Triclinic  system  (Fig.  50).    Three  axes,   all  inclined  to  each 
other. 

6.  Hexagonal  system  (Fig.  51).    Three  equal  axes  in  the  same 
plane  intersecting  at  angles  of  60°,  and  a  fourth  at  right  angles  to 
all  of  these. 

Every  crystal  can  be  imagined  to  have  its  faces  and  angles 
arranged  in  a  definite  way  around  one  of  these  systems  of  axes.  A 
cube,  for  instance,  is  referred  to  Plan  i,  an  axis  ending  in  the  center 
of  each  face ;  while  in  a  regular  octohedron  an  axis  ends  in  each 
solid  angle.  These  forms  are  shown  in  Fig.  46.  It  will  be  seen  that 
both  of  these  figures  belong  to  the  same  system,  though  they  are 
very  different  in  appearance.  In  the  same  way,  many  geometric 


SULPHUR  AND  ITS  COMPOUNDS  163 

forms  may  be  derived  from  each  of  the  systems,  and  the  light  lines 
about  the  axes  in  the  drawings  show  two  of  the  simplest  forms  of 
each  of  the  systems. 

In  general  a  given  substance  always  crystallizes  in  the  same  sys- 
tem, and  two  corresponding  faces  of  each  crystal  of  it  always  make 
the  same  angle  with  each  other.  A  few  substances,  of  which  sulphur 
is  an  example,  crystallize  in  two  different  systems,  and  the  crystals 
differ  in  such  physical  properties  as  melting  point  and  density.  Such 
substances  are  said  to  be  dimorphous. 


EXERCISES 

1.  (#)  Would  the  same  amount  of  heat  be  generated  by  the  com- 
bustion of  i  g.  of  each  of  the  allotropic  modifications  of  sulphur  ? 
(£)  Would  the  same  amount  of  sulphur  dioxide  be  formed  in  each 
case  ? 

2.  Is  the  equation  for  the  preparation  of  hydrosulphuric  acid  a 
reversible  one?    As  ordinarily  carried  out,  does  the  reaction  com- 
plete itself? 

3.  Suppose  that  hydrosulphuric  acid  were  a  liquid,  would  it  be 
necessary  to  modify  the  method  of  preparation  ? 

4.  Can  sulphuric  acid  be  used  to  dry  hydrosulphuric  acid  ?    Give 
reason  for  answer. 

5.  Does  dry  hydrosulphuric  acid  react  with  litmus  paper  ?    State 
reason  for  answer. 

6.  How  many  grams  of  iron  sulphide  are  necessary  to  prepare 
100  1.  of  hydrosulphuric  acid  when   the  laboratory  conditions  are 
17°  and  740  mm.  pressure? 

7.  Suppose  that  the  hydrogen  in  i  1.  of  hydrosulphuric  acid  were 
liberated  ;  what  volume  would  it  occupy,  the  gases  being  measured 
under  the  same  conditions  ? 

8.  Write  the  equations  representing  the  reaction  between  hydro- 
sulphuric    acid    and   sodium   hydroxide    and   ammonium   hydroxide 
respectively. 

9.  Show  that  the  preparation  of  sulphur  dioxide  from  a  sulphite 
is  similar  in  principle  to  the  preparation  of  hydrogen  sulphide. 

10.  (a)  Does  dry  sulphur  dioxide  react  with  litmus  paper? 
(£)  How  can  it  be  shown  that  a  solution  of  sulphur  dioxide  in  water 
acts  like  an  acid  ? 


1 64     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

11.  (a)  Calculate  the  percentage  composition  of  sulphurous  anhy- 
dride and  sulphuric  anhydride.    (£)  Show  how  these  two  substances 
are  in  harmony  with  the  law  of  multiple  proportion. 

12.  How  many  pounds  of  sulphur  would  be  necessary  in  the  prep- 
aration of  100  Ib.  of  98%  sulphuric  acid? 

13.  What  weight  of  sulphur  dioxide  is  necessary  in  the  prepara- 
tion of  i  kg.  of  sodium  sulphite  ? 

14.  What  weight  of  copper  sulphate  crystals  can  be  obtained  by 
dissolving   i  kg.  of  copper  in   sulphuric   acid  and  crystallizing  the 
product  from  water  ? 

15.  Write  the  names  and  formulas  of  the  oxides  and  oxygen  acids 
of  selenium  and  tellurium. 

16.  In  the  commercial  preparation  of  carbon  disulphide,  what  is 
the  function  of  the  electric  current  ? 

17.  If  the  Gay-Lussac  tower  were  omitted  from  the  sulphuric  acid 
factory,  what  effect  would  this  have  on  the  cost  of  production  of 
sulphuric  acid? 


CHAPTER   XV 
PERIODIC  LAW 

A  number  of  the  elements  have  now  been  studied  some- 
what closely.  The  first  three  of  these,  oxygen,  hydrogen, 
and  nitrogen,  while  having  some  physical  properties  in 
common  with  each  other,  have  almost  no  point  of  similarity 
as  regards  their  chemical  conduct.  On  the  other  hand, 
oxygen  and  sulphur,  while  quite  different  physically,  have 
much  in  common  in  their  chemical  properties. 

About  eighty  elements  are  now  known.  If  all  of  these 
should  have  properties  as  diverse  as  do  oxygen,  hydrogen, 
and  nitrogen,  the  study  of  chemistry  would  plainly  be  a  very 
difficult  and  complicated  one.  If,  however,  the  elements 
can  be  classified  in  groups,  the  members  of  which  have  very 
similar  properties,  the  study  will  be  very  much  simplified. 

Earlier  classification  of  the  elements.  Even  at  an  early 
period  efforts  were  made  to  discover  some  natural  principle 
in  accordance  with  which  the  elements  could  be  classified. 
Two  of  these  classifications  may  be  mentioned  here. 

i .  Classification  into  metals  and  non-metals.  The  classifi- 
cation into  metals  and  non-metals  most  naturally  suggested 
itself.  This  grouping  was  based  largely  on  physical  prop- 
erties, the  metals  being  heavy,  lustrous,  malleable,  ductile, 
and  good  conductors  of  heat  and  electricity.  Elements 
possessing  these  properties  are  usually  base-forming  in 
character,  and  the  ability  to  form  bases  came  to  be 
regarded  as  a  characteristic  property  of  the  metals.  The 

165 


1 66     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

non-metals  possessed  physical  properties  which  were  the 
reverse  of  those  of  the  metals,  and  were  acid-forming  in 
character. 

Not  much  was  gained  by  this  classification,  and  it  was 
very  imperfect.  Some  metals,  such  as  potassium,  are  very 
light ;  some  non-metals,  such  as  iodine,  have  a  high  luster ; 
some  elements  can  form  either  an  acid  or  a  base. 

2.  Classification  into  triad  families.  In  1825  Dobereiner 
observed  that  an  interesting  relation  exists  between  the 
atomic  weights  qf  chemically  similar  elements.  To  illus- 
trate, lithium,  sodium,  and  potassium  resemble  each  other 
very  closely,  and  the  atomic  weight  of  sodium  is  almost 
exactly  an  arithmetical  mean  between  those  of  the  other 

7-°3  +  SQ^S 
two  :  -  -  =  23.09.    In  many  chemical  and  physi- 

2 

cal  properties  sodium  is  midway  between  the  other  two. 

A  number  of  triad  families  were  found,  but  among  eighty 
elements,  whose  atomic  weights  range  all  the. way  from 
i  to  240,  such  agreements  might  be  mere  chance.  Moreover 
many  elements  did  not  appear  to  belong  to  such  families. 

Periodic  division.  In  1869  the  Russian  chemist  Men- 
deleeff  devised  an  arrangement  of  the  elements  based 
on  their  atomic  weights,  which  has  proved  to  be  of  great 
service  in  the  comparative  study  of  the  elements.  A  few 
months  later  the  German,  Lothar  Meyer,  independently 
suggested  the  same  ideas.  This  arrangement  brought  to 
light  a  great  generalization,  now  known  as  the  periodic  law. 
An  exact  statement  of  the  law  will  be  given  after  the  method 
of  arranging  the  elements  has  been  described. 

Arrangement  of  the  periodic  table.  The  arrangement 
suggested  by  Mendeleeff,  modified  somewhat  by  more  re- 
cent investigations,  is  as  follows  :  Beginning  with  lithium, 


PERIODIC  LAW  167 

which  has  an  atomic  weight  of  7,  the  elements  are  arranged 
in  a  horizontal  row  in  the  order  of  their  atomic  weights, 
thus  : 

Li  (7.03),  Be  (9.1),  B  (n),  C  (12),  N  (14.04),  0  (16),  F  (19). 

These  seven  elements  all  differ  markedly  from  each  other. 
The  eighth  element,  sodium,  is  very  similar  to  lithium.  It 
is  placed  just  under  lithium,  and  a  new  row  follows  : 

Na  (23.05),  Mg  (24.36),  Al  (27.  i),  Si  (28.4),  P (31),  8(32.06),  Cl (35.45). 

When  the  fifteenth  element,  potassium,  is  reached,  it  is 
placed  under  sodium,  to  which  it  is  very  similar,  and  serves 
to  begin  a  third  row : 

K  (39-15),  Ca  (40.1),  Sc  (44.1,)  Ti  (48.1),  V  (51.2), Cr  (52.1),  Mn(55). 

Not  only  is  there  a  strong  similarity  between  lithium,  so- 
dium, and  potassium,  which  have  been  placed  in  a  vertical 
row  because  of  this  resemblance,  but  the  elements  in  the 
other  vertical  rows  exhibit  much  of  the  same  kind  of  simi- 
larity among  themselves,  and  evidently  form  little  natural 
groups. 

The  three  elements  following  manganese,  namely, 'iron, 
nickel,  and  cobalt,  have  atomic  weights  near  together,  and 
are  very  similar  chemically.  They  do  not  strongly  resemble 
any  of  the  elements  so  far  considered,  and  are  accordingly 
placed  in  a  group  by  themselves,  following  manganese.  A 
new  row  is  begun  with  copper,  which  somewhat  resembles 
the  elements  of  the  first  vertical  column.  Following  the 
fifth  and  seventh  rows  are  groups  of  three  closely  related 
elements,  so  that  the  completed  arrangement  has  the  ap- 
pearance represented  in  the  table  on  page  168. 

Place  of  the  atmospheric  elements.  When  argon  was 
discovered  it  was  seen  at  once  that  there  was  no  place  in 


II     II 


II  II 


da 


l> 


OK 


OS 


& 


3S 


168 


PERIODIC  LAW  169 

the  table  for  an  element  of  atomic  weight  approximately  40. 
When  the  other  inactive  elements  were  found,  however,  it 
became  apparent  that  they  form  a  group  just  preceding 
Group  I.  They  are  accordingly  arranged  in  this  way  in 
Group  0  (see  table  on  opposite  page).  A  study  of  this 
table  brings  to  light  certain  very  striking  facts. 

Properties  of  elements  vary  with  atomic  weights.  There 
is  evidently  a  close  relation  between  the  properties  of  an 
element  and  its  atomic  weight.  Lithium,  at  the  beginning 
of  the  first  group,  is  a  very  strong  base-forming  element, 
with  pronounced  metallic  properties.  Beryllium,  following 
lithium,  is  less  strongly  base-forming,  while  boron  has  some 
base-forming  and  some  acid-forming  properties.  In  carbon 
all  base-formjng  properties  have  disappeared,  and  the  acid- 
forming  properties  are  more  marked  than  in  boron.  These 
become  still  more  emphasized  as  we  pass  through  nitrogen 
and  oxygen,  until  on  reaching  fluorine  we  have  one  of  the 
strongest  acid-forming  elements.  The  properties  of  these 
seven  elements  therefore  vary  regularly  with  their  atomic 
weights,  or,  in  mathematical  language,  are  regular  functions 
of  them. 

Periodic  law.  The  properties  of  the  first  seven  elements 
vary  continuously  —  that  is  steadily  —  away  from  base- 
forming  and  toward  acid-forming  properties.  If  lithium  had 
the  smallest  atomic  weight  of  any  of  the  elements,  and  fluor- 
ine the  greatest,  so  that  in  passing  from  one  to  the  other  we 
had  included  all  the  elements,  we  could  say  that  the  properties 
of  elements  are  continuous  functions  of  their  atomic  weights. 
But  fluorine  is  an  element  of  small  atomic  weight,  and  the 
one  following  it,  sodium,  breaks  the  regular  order,  for  in  it 
reappear  all  the  characteristic  properties  of  lithium.  Mag- 
nesium, following  sodium,  bears  much  the  same  relation  to 


170     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

beryllium  that  sodium  does  to  lithium,  and  the  properties 
of  the  elements  in  the  second  row  vary  much  as  they  do 
in  the  first  row  until  potassium  is  reached,  when  another 
repetition  begins.  The  properties  of  the  elements  do  not 
vary  continuously,  therefore,  with  atomic  weights,  but  at 
regular  intervals  there  is  a  repetition,  or  period.  This  gen- 
eralization is  known  as  the  periodic  /aw,  and  may  be  stated 
thus  :  The  properties  of  elements  are  periodic  functions  of 
their  atomic  weights. 

The  two  families  in  a  group.  While  all  the  elements  in 
a  given  vertical  column  bear  a  general  resemblance  to  each 
other,  it  has  been  noticed  that  those  belonging  to  periods 
having  even  numbers  are  very  strikingly  similar  to  each 
other.  They  are  placed  at  the  left  side  of  the  group  col- 
umns. In  like  manner,  the  elements  belonging  to  the  odd 
periods  are  very  similar  and  are  arranged  at  the  right  side 
of  the  group  columns.  Thus  calcium,  strontium,  and  barium 
are  very  much  alike  ;  so,  too,  are  magnesium,  zinc,  and  cad- 
mium. The  resemblance  between  calcium  and  magnesium, 
or  strontium  and  zinc,  is  much  less  marked.  This  method 
of  arrangement  therefore  divides  each  group  into  two  fam- 
ilies, each  containing  four  or  five  members,  between  which 
there  is  a  great  similarity. 

Family  resemblances.  Let  us  now  inquire  more  closely 
in  what  respects  the  elements  of  a  family  resemble  each 
other. 

i.  Valence.  In  general  the  valence  of  the  elements  in  a 
family  is  the  same,  and  the  formulas  of  their  compounds 
are  therefore  similar.  If  we  know  that  the  formula  of  so- 
dium chloride  is  NaCl,  it  is  pretty  certain  that  the  formula 
of  potassium  chloride  will  be  KC1  —  not  KC12  or  KC13. 
The  general  formulas  R2O,  RO,  etc.,  placed  below  the 


PERIODIC  LAW  171 

columns  show  the  formulas  of  the  oxides  of  the  elements 
in  the  column  provided  they  form  oxides.  In  like  manner 
the  formulas  RH,  RH2,  etc.,  show  the  composition  of  the 
compounds  formed  with  hydrogen  or  chlorine. 

2.  Chemical  properties.    The  chemical  properties  of  the 
members  of  a  family  are  quite  'similar.    If  one  member  is 
a  metal,  the  others  usually  are ;  if  one  is  a  non-metal,  so, 
too,  are  the  others.    The  families  in  the  first  two  columns 
consist  of  metals,  while  the  elements  found  in  the  last  two 
columns  form  acids.    There  is  in  addition  a  certain  regu- 
larity in  properties  of  the  elements  in  each  family.    If  the 
element  at  the  head  of  the  family  is  a  strong  acid-forming 
element,  this  property  is  likely  to  diminish  gradually,  as 
we  pass  to  the  members  of  the  family  with  higher  atomic 
weights.    Thus  phosphorus  is  strongly  acid-forming,  arsenic 
less' so,  antimony  still  less  so,  while  bismuth  has  almost  no 
acid-forming  properties.    We  shall  meet  with  many  illus- 
trations of  this  fact. 

3.  Physical  properties.    In  the  same  way,  the  physical 
properties  of  the  members  of  a  family  are  in  general  some- 
what similar,  and  show  a  regular  gradation  as  we  pass  from 
element  to  element  in  the  family.    Thus  the  densities  of 
the  members  of  the  magnesium  family  are 

Mg  =  1.75,  Zn  =  7.00,  Cd  =  8.6;,  Hg  =  13.6. 
Their  melting  points  are 

Mg  =  750°,  Zn  =  420°,  Cd  =  320°,  Hg  =  -39-5°. 
Value  of  the  periodic  law.    The  periodic  law  has  proved 
of   much   value   in   the   development   of    the   science   of 
chemistry. 

i.  It  simplifies  study.    It  is  at  once  evident  that  such 
regularities  very  much  simplify  the  study  of  chemistry. 


172     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

A  thorough  study  of  one  element  of  a  family  makes  the 
study  of  the  other  members  a  much  easier  task,  since  so 
many  of  the  properties  and  chemical  reactions  of  the  ele- 
ments are  similar.  Thus,  having  studied  the  element  sul- 
phur in  some  detail,  it  is  not  necessary  to  study  selenium 
and  tellurium  so  closely,  for  most  of  their  properties  can 
be  predicted  from  the  relation  which  they  sustain  to 
sulphur. 

2.  It  predicts  new  elements.    When  the  periodic  law  was 
first  formulated  there  were  a  number  of  vacant  places  in 
the  table  which   evidently  belonged  to  elements  at  that 
time  unknown.    From   their  position  in  the   table,   Men- 
deleeff  predicted   with  great  precision  the  properties   of 
the  elements  which  he  felt  sure  would  one  day  be  discov- 
ered to  fill  these  places.    Three  of  them,  scandium,  ger- 
manium, and  gallium,  were  found  within  fifteen  years,  and 
their  properties  agreed  in  a  remarkable  way  with  the  pre- 
dictions of  Mendele"eff .    There  are  still  some  vacant  places 
in  the  table,  especially  among  the  heavier  elements. 

3.  It  corrects  errors.    The  physical  constants  of  many  of 
the  elements  did  not  at  first  agree  with  those  demanded 
by  the  periodic  law,  and  a  further  study  of  many  such 
cases  showed  that  errors  had  been  made.    The  law  has 
therefore  done  much  service  in  indicating  probable  error. 

Imperfections  of  the  law.  There  still  remain  a  good 
many  features  which  must  be  regarded  as  imperfections  in 
the  law.  Most  conspicuous  is  the  fact  that  the  element 
hydrogen  has  no  place  in  the  table.  In  some  of  the  groups 
elements  appear  in  one  of  the  families,  while  all  of  their 
properties  show  that  they  belong  in  the  other.  Thus 
sodium  belongs  with  lithium  and  not  with  copper ;  fluorine 
belongs  with  chlorine  and  not  with  manganese.  There  are 


PERIODIC  LAW  173 

two  instances  where  the  elements  must  be  transposed  in 
order  to  make  them  fit  into  their  proper  group.  According 
to  their  atomic  weights,  tellurium  should  follow  iodine,  and 
argon  should  follow  potassium.  Their  properties  show  in 
each  case  that  this  order  must  be  reversed.  The  table 
separates  some  elements  altogether  which  in  many  respects 
have  closely  agreeing  properties.  Iron,  chromium,  and  man- 
ganese are  all  in  different  groups,  although  they  are  similar 
in  many  respects. 

The  system  is  therefore  to  be  regarded  as  but  a  partial 
and  imperfect  expression  of  some  very  important  and  fun- 
damental relation  between  the  substances  which  we  know 
as  elements,  the  exact  nature  of  this  relation  being  as  yet 
not  completely  clear  to  us. 

EXERCISES 

1.  Suppose  that  an  element  were  discovered  that  filled  the  blank 
in  Group  0,  Period  5 ;  what  properties  would  it  probably  have  ? 

2.  Suppose  that  an  element  were  discovered  that  filled  the  blank 
in  Group  VI,  Period  9,  family  B',  what  properties  would  it  have? 

3.  Sulphur  and  oxygen  both  belong  in  Group  VI,  although  in  dif- 
ferent families ;  in  what  respects  are  the  two  similar? 


CHAPTER  XVI 
THE  CHLORINE  FAMILY 


ATOMIC 
WEIGHT 

MELTING 
POINT 

BOILING 
POINT 

COLOR  AND  STATE 

Fluorine  (F)   .     . 

19.00 

-223° 

-187° 

Pale  yellowish  gas. 

Chlorine  (Cl)  .     . 

35-45 

-102° 

-33-6° 

Greenish-yellow  gas. 

Bromine  (Br)  .     . 

79.96 

-7° 

59° 

Red  liquid. 

Iodine  (I)  .     .     . 

126.97 

107° 

i?5° 

Purplish-black  solid. 

The  family.  The  four  elements  named  in  the  above 
table  form  a  strongly  marked  family  of  elements  and  illus- 
trate very  clearly  the  way  in  which  the  members  of  a  fam- 
ily in  a  periodic  group  resemble  each  other,  as  well  as  the 
character  of  the  differences  which  we  may  expect  to  find 
between  the  individual  members. 

1.  Occurrence.    These  elerrjents  do  not  occur  in  nature 
in  the  free  state.    The  compounds  of  the  last  three  ele- 
ments of  the  family  are  found  extensively  in  sea  water, 
and  on  this  account  the  name  halogens,  signifying  "  pro- 
ducers of  sea  salt,"  is  sometimes  applied  to  the  family. 

2.  Properties.    As  will  be  seen  by  reference  to  the  table, 
the  melting  points  and  boiling  points  of  the  elements  of 
the  family  increase  with  their  atomic  weights.     A  some- 
what  similar  gradation   is  noted  in  their  color  and  state. 
One  atom  of  each  of  the  elements  combines  with  one  atom 
of  hydrogen  to  form  acids,  which  are  gases  very  soluble 
in  water.    The  affinity  of  the  elements  for  hydrogen  is  in 

174 


THE  CHLORINE  FAMILY 


175 


the  inverse  order  of  their  atomic  weights,  fluorine  hav- 
ing the  strongest  affinity  and  iodine  the  weakest.  Only 
chlorine  and  iodine  form  oxides,  and  those  of  the  former 
element  are  very  unstable.  The  elements  of  the  group 
are  univalent  in  their  compounds  with  hydrogen  and  the 
metals. 

FLUORINE 

Occurrence.  The  element  fluorine  occurs  in  nature 
most  abundantly  as  the  mineral  fluorspar  (CaF2),  as  cry- 
olite (Na3AlF6),  and  in  the  complex  mineral  apatite 
(3Ca3(P04)2-CaF2). 

Preparation.  All  attempts  to  isolate  the  element  resulted 
in  failure  until  recent  years.  Methods  similar  to  those  which 
succeed  in  the  preparation  of  the  other  elements  of  the 
family  cannot  be  used  ;  for  as  soon  as 
the  fluorine  is  liberated  it  combines 
with  the  materials  of  which  the  appa- 
ratus is  made  or  with  the  hydrogen  of 
the  water  which  is  always  present.  g 
The  preparation  of  fluorine  was  finally 
accomplished  by  the  French  chemist 
Moissan  by  the  electrolysis  of  hydro- 
fluoric acid.  Perfectly  dry  hydrofluoric 
acid  (HF)  was  condensed  to  a  liquid 
and  placed  in  a  U-shaped  tube  made 
of  platinum  (or  copper),  which  was 
furnished  with  electrodes  and  delivery  tubes,  as  shown  in 
Fig.  52.  This  liquid  is  not  an  electrolyte,  but  becomes 
such  when  potassium  fluoride  is  dissolved  in  it.  When 
this  solution  was  electrolyzed  hydrogen  was  set  free  at 
the  cathode  and  fluorine  at  the  anode. 


FIG.  52 


176     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Properties.  Fluorine  is  a  gas  of  slightly  yellowish  color, 
and  can  be  condensed  to  a  liquid  boiling  at  —  187°  under 
atmospheric  pressure.  It  solidifies  at  —223°.  It  is  ex- 
tremely active  chemically,  being  the  most  active  of  all 
the  elements  at  ordinary  temperatures. 

It  combines  with  all  the  common  elements  save  oxygen, 
very  often  with  incandescence  and  the  liberation  of  much 
heat.  It  has  a  strong  affinity  for  hydrogen  and  is  able 
to  withdraw  it  from  its  compounds  with  other  elements. 
Because  of  its  great  activity  it  is  extremely  poisonous. 
Fluorine  does  not  form  any  oxides,  neither  does  it  form 
any  oxygen  acids,  in  which  respects  it  differs  from  the 
other  members  of  the  family. 

Hydrofluoric  acid  (HF).  Hydrofluoric  acid  is  readily 
obtained  from  fluorspar  by  the  action  of  concentrated 
sulphuric  acid.  The  equation  is 

CaF2  +  H2SO4  =.CaSO4  +  2  HF. 

In  its  physical  properties  it  resembles  the  binary  acids  of 
the  other  elements  of  this  family,  being,  however,  more 
easily  condensed  to  a  liquid.  The  anhydrous  acid  boils  at 
19°  and  can  therefore  be  prepared  at  ordinary  pressures. 
It  is  soluble  in  all  proportions  in  water,  and  a  concentrated 
solution  —  about  50/0  —  is  prepared  for  the  market.  Its 
fumes  are  exceedingly  irritating  to  the  respiratory  organs, 
and  several  chemists  have  lost  their  lives  by  accidentally 
breathing  them. 

Chemical  properties.  Hydrofluoric  acid,  like  other  strong 
acids,  readily  acts  on  bases  and  metallic  oxides  and  forms 
the  corresponding  fluorides.  It  also  dissolves  certain  metals 
such  as  silver  and  copper.  It  acts  very  vigorously  upon 
organic  matter,  a  single  drop  of  the  concentrated  acid 


THE  CHLORINE  FAMILY  177 

making  a  sore  on  the  skin  which  is  very  painful  and  slow 
in  healing.  Its  most  characteristic  property  is  its  action 
upon  silicon  dioxide  (SiO2),  with  which  it  forms  water  and 
the  gas  silicon  tetrafluoride  (SiF4),  as  shown  in  the  equation 

SiO2  +  4  HF  =  SiF4  +  2  H2O. 

Glass  consists  of  certain  compounds  of  silicon,  which  are 
likewise  acted  on  by  the  acid  so  that  it  cannot  be  kept 
in  glass  bottles.  It  is  preserved  in  flasks  made  of  wax  or 
gutta-percha. 

Etching.  Advantage  is  taken  of  this  reaction  in  etching  designs 
upon  glass.  The  glass  vessel  is  painted  over  with  a  protective  paint 
upon  which  the  acid  will  not  act,  the  parts  which  it  is  desired  to 
make  opaque  being  left  unprotected.  A  mixture  of  fluorspar  and 
sulphuric  acid  is  then  painted  over  the  vessel  and  after  a  few  min- 
utes the  vessel  is  washed  clean.  Wherever  the,  hydrofluoric  acid 
comes  in  contact  with  the  glass  it  acts  upon  it,  destroying  its  luster 
and  making  it  opaque,  so  that  the  exposed  design  will  be  etched 
upon  the  clear  glass.  Frosted  glass  globes  are  often  made  in  this 
way. 

The  etching  may  also  be  effected  by  covering  the  glass  with  a 
thin  layer  of  paraffin,  cutting  the  design  through  the  wax  and  then 
exposing  the  glass  to  the  fumes  of  the  acid. 

Salts  of  hydrofluoric  acid, — fluorides.  A  number  of 
the  fluorides  are  known,  but  only  one  of  them,  calcium 
fluoride  (CaF2),  is  of  importance.  This,  is  the  well-known 
mineral  fluorspar. 

CHLORINE 

Historical.  While  studying  the  action  of  hydrochloric 
acid  upon  the  mineral  pyrolusite,  in  1774,  Scheele  obtained 
a  yellowish,  gaseous  substance  to  which  he  gave  a  name  in 
keeping  with  the  phlogiston  theory  then  current.  Later  it 
was  supposed  to  be  a  compound  containing  oxygen.  In 


1 78     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

1810,  however,  the  English  chemist  Sir  Humphry  Davy 
proved  it  to  be  an  element  and  named  it  chlorine. 

Occurrence.  Chlorine  does  not  occur  free  in  nature,  but 
its  compounds  are  widely  distributed.  For  the  most  part 
it  occurs  in  combination  with  the  metals  in  the  form  of 
chlorides,  those  of  sodium,  potassium,  and  magnesium 
being  most  abundant.  Nearly  all  salt  water  contains  these 
substances,  particularly  sodium  chloride,  and  very  large 
salt  beds  consisting  of  chlorides  are  found  in  many  parts 
of  the  world. 

Preparation.  Two  general  methods  of  preparing  chlo- 
rine may  be  mentioned,  namely,  the  laboratory  method  and 
the  electrolytic  method. 

i .  Laboratory  method.  In  the  laboratory  chlorine  is  made 
by  warming  the  mineral  pyrolusite  (manganese  dioxide, 
MnO2)  with  concentrated  hydrochloric  acid.  The  first  re- 
action, which  seems  to  be  similar  to  the  action  of  acids 
upon  oxides  in  general,  is  expressed  in»the  equation 

MnO2  +  4  HC1  -  MnCl4  +  2  H2O. 

The  manganese  compound  so  formed  is  very  unstable, 
however,  and  breaks  down  according  to  the  equation 

MnCl4  =  MnCl2  +  2  Cl. 

Instead  of  using  hydrochloric  acid  in  the  preparation  of 
chlorine  it  will  serve  just  as  well  to  use  a  mixture  of  so- 
dium chloride  and  sulphuric  acid,  since  these  two  react  to 
form  hydrochloric  acid.  The  following  equations  will  then 
express  the  changes : 

(1 )  2  NaCl  +  H2SO4  =  Na2SO4  +  2  HC1 . 

(2)  MnO2  +  4  HC1  =  MnCl2  +  2  Cl  +  2  H2O. 

(3)  MnCfa  +  H2SO4  =  MnSO4  +  2  HC1. 


THE  CHLORINE  FAMILY 


179 


Combining  these  equations,  the  following  equation  express- 
ing the  complete  reaction  is  obtained  : 

2  NaCl  +  MnO2  +  2  H2SO4  =  MnSO4  +  Na2SO4 


Since  the  hydrochloric  acid  liberated  in  the  third  equation 
is  free  to  act  upon  manganese  dioxide,  it  will  be  seen  that 

all  of    the  chlorine  originally  present  in    the 

sodium  chloride  is  set  free. 

The  manganese  dioxide  and  the  hydrochloric  acid 
are  .brought  together  in  a  flask,  as  repre- 
sented in  Fig.  53,  and  a  gentle  heat  is  ap- 
plied. The  rate  of  evolution  of  the  gas  is 
regulated  by  the  amount  of  heat  applied, 
and  the  gas  is  collected  by  displacement  of 
air.  As  the  equations  show,  only  half  of 
the  chlorine  present  in  the  hydrochloric 
acid  is  liberated. 


2.  Electrolytic  method. 
Under  the  discussion  of  elec- 
trolysis (p.  102)  it  was  shown 
that  when  a  solution  of  sodium 
chloride  is  electrolyzed  chlorine 


FIG. 53 


is  evolved  at  the  anode,  while  the  sodium  set  free  at  the 
cathode  reacts  with  the  water  to  form  hydrogen,  which  is 
evolved,  and  sodium  hydroxide,  which  remains  in  solution. 
A  great  deal  of  the  chlorine  required  in  the  chemical  in- 
dustries is  now  made  in  this  way  in  connection  with  the 
manufacture  of  sodium  hydroxide. 

Physical  properties.  Chlorine  is  a  greenish-yellow  gas 
which  has  a  peculiar  suffocating  odor  and  produces  a  very 
violent  effect  upon  the  throat  and  lungs.  Even  when  inhaled 
in  small  quantities  it  often  produces  all  the  symptoms  of  a 


l8o    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

hard  cold,  and  in  larger  quantities  may  have  serious  and 
even  fatal  action.  It  is  quite  heavy  (density  =  2.45).  and 
can  therefore  be  collected  by  displacement  of  air.  One 
volume  of  water  under  ordinary  conditions  dissolves  about 
three  volumes  of  chlorine.  The  gas  is  readily  liquefied,  a 
pressure  of  six  atmospheres  serving  to  liquefy  it  at  o°.  It 
forms  a  yellowish  liquid  which  solidifies  at  —  102°. 

Chemical  properties.  At  ordinary  temperatures  chlorine 
is  far  more  active  chemically  than  any  of  the  elements 
we  have  so  far  considered,  with  the  exception  of  fluorine  ; 
indeed,  it  is  one  of  the  most  active  of  all  elements. 

1.  Action    on    metals.    A   great   many   metals    combine 
directly  with   chlorine,   especially  when  hot.    A   strip   of 
copper  foil  heated  in  a  burner  flame  and  then  dropped 
into   chlorine   burns   with   incandescence.     Sodium   burns 
brilliantly  when  heated  strongly  in  slightly  moist  chlorine. 
Gold  and  silver  are  quickly  tarnished  by  the  gas. 

2.  Action  on  non-metals.    Chlorine  has  likewise  a  strong 
affinity  for    many    of    the  non-metals.     Thus    phosphorus 
burns  in  a  current  of  the  gas,  while  antimony  and  arsenic 
in  the  form  of  a  fine  powder  at  once  burst  into  flame  when 
dropped  into  jars  of  the  gas.    The  products  formed  in  all 
cases  where  chlorine  combines  with  another  element  are 
called  chlorides. 

3.  Action  on  hydrogen.    Chlorine  has  a  strong  affinity 
for  hydrogen,  uniting  with  it  to  form  hydrochloric  acid. 
A  jet  of  hydrogen  burning  in  the  air  continues  to  burn 
when  introduced  into  a  jar  of  chlorine,  giving  a  somewhat 
luminous  flame.    A  mixture  of  the  two  gases  explodes  vio- 
lently when  a  spark  is  passed   through   it  or  when  it  is 
exposed  to  bright   sunlight.    In  the  latter  case  it  is  the 
light  and  not  the  heat  which  starts  the  action. 


THE  CHLORINE  FAMILY  l8l 

4.  Action  on  substances  containing  hydrogen.  Not  only 
will  chlorine  combine  directly  with  free  hydrogen  but  it 
will  often  abstract  the  element  from  its  compounds.  Thus, 
when  chlorine  is  passed  into  a  solution  containing  hydro- 
sulphuric  acid,  sulphur  is  precipitated  and  hydrochloric 
acid  formed.  The  reaction  is  shown  by  the  following 
equation  :  +  2  Cl  =  2  HC1  +  S. 


With  ammonia  the  action  is  similar  : 


The  same  tendency  is  very  strikingly  seen  in  the  action 
of  chlorine  upon  turpentine.  The  latter  substance  is  largely 
made  up  of  compounds  having  the  composition 
represented  by  the  formula  C10H16.  When  a 
strip  of  paper  moistened  with  warm  turpentine 
is  placed  in  a  jar  of  chlorine  dense  fumes  of 
hydrochloric  acid  appear  and  a  black  deposit  of 
carbon  is  formed.  Even  water,  which  is  a  very 
stable  compound,  can  be  decomposed  by  chlorine, 
the  oxygen  being  liberated.  This  may  be  shown 
in  the  following  way  : 

If  a  long  tube  of  rather  large  diameter  is  filled  with  a 
strong  solution  of  chlorine  in  water  and  inverted  in  a 
vessel  of  the  same  solution,  as  shown  in  Fig.  54,  and  the 
apparatus  is  placed  in  bright  sunlight,  very  soon  bubbles 
of  a  gas  will  be  observed  to  rise  through  the  solution  and  collect  in 
the  tube.    An  examination  of  this  gas  will  show  that  it  is  oxygen.    It 
is  liberated  from  water  in  accordance  with  the  following  equation : 

H20  +  2  Cl  =  2  HC1  +  O. 

5.  Action  on  color  substances,  —  bleaching  action.  If 
strips  of  brightly  colored  cloth  or  some  highly  colored 
flowers  are  placed  in  quite  dry  chlorine,  no  marked  change 


182     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

in  color  is  noticed  as  a  rule.  If,  however,  the  cloth  and 
flowers  are  first  moistened,  the  color  rapidly  disappears, 
that  is,  the  objects  are  bleached.  Evidently  the  moisture 
as  well  as  the  chlorine  is  concerned  in  the  action,  and  a 
study  of  the  case  shows  that  the  chlorine  has  combined 
with  the  hydrogen  of  the  water.  The  oxygen  set  free 
oxidizes  the  color  substance,  converting  it  into  a  colorless 
compound.  It  is  evident  from  this  explanation  that  chlorine 
will  only  bleach  those  substances  which  are  changed  into 
colorless  compounds  by  oxidation. 

6.  Action  as  a  disinfectant.  Chlorine  has  also  marked 
germicidal  properties,  and  the  free  element,  as  well  as 
compounds  from  which  it  is  easily  liberated,  are  used  as 
disinfectants. 

Nascent  state.  It  will  be  noticed  that  oxygen  when  set 
free  from  water  by  chlorine  is  able  to  do  what  ordinary 
oxygen  cannot  do,  for  both  the  cloth  and  the  flowers  are 
unchanged  in  the  air  which  contains  oxygen.  It  is  generally 
true  that  the  activity  of  an  element  is  greatest  at  the  instant 
of  liberation  from  its  compounds.  To  express  this  fact 
elements  at  the  instant  of  liberation  are  said  to  be  in  the 
nascent  state.  It  is  nascent  oxygen  which  does  the  bleaching. 

Hydrochloric  acid  (muriatic  acid)  (HC1).  The  preparation 
of  hydrochloric  acid  may  be'  discussed  under  two  general 
heads : 

i.  Laboratory  preparation.  The  product  formed  by  the 
burning  of  hydrogen  in  chlorine  is  the  gas  hydrochloric 
acid.  This  substance  is  much  more  easily  obtained,  how- 
ever, by  treating  common  salt  (sodium  chloride)  with 
sulphuric  acid/  The  following  equation  shows  the  reaction  : 

2  NaCl  +  H2SO4  =  Na2SO4  +  2  HC1. 


THE  CHLORINE  FAMILY  183 

The  dry  salt  is  placed  in  a  flask  furnished  with  a  funnel 
tube  and  an  exit,  tube,  the  sulphuric  acid  is  added,  and  the 
flask  gently  warmed.  The  hydrochloric  acid  gas  is  rapidly 
given  off  and  can  be  collected  by  displacement  of  air.  The 
same  apparatus  can  be  used  as  was  employed  in  the  prepa- 
ration of  chlorine  (Fig.  53). 

When  a  solution  of  salt  is  treated  with  sulphuric  acid  there  is  no 
very  marked  action.  The  hydrochloric  acid  formed  is  very  soluble 
in  water,  and  so  does  not  escape  from  the  solution ;  hence  a  state  of 
equilibrium  is  soon  reached  between  the  four  substances  represented 
in  the  equation.  When  concentrated  sulphuric  acid,  in  which  hydro- 
chloric acid  is  not  soluble,  is  poured  upon  dry  salt  the  reaction  is 
complete. 

2.  Commercial  preparation.  Commercially,  hydrochloric 
acid  is  prepared  in  connection  with  the  manufacture  of 
sodium  sulphate,  the  reaction  being  the  same  as  that  just 
given.  The  reaction  is  carried  out  in  a  furnace,  and  the 
hydrochloric  acid  as  it  escapes  in  the  form  of  gas  is  passed 
into  water  in  which  it  dissolves,  the  solution  forming  the 
hydrochloric  acid  of  commerce.  When  the  materials  are 
pure  a  colorless  solution  is  obtained.  The  most  concen- 
trated solution  has  a  density  of  1.2  and  contains  40%  HC1. 
The  commercial  acid,  often  called  muriatic  acid,  is  usually 
colored  yellow  by  impurities. 

Composition  of  hydrochloric  acid.  When  a  solution  of 
hydrochloric  acid  is  electrolyzed  in  an  apparatus  similar  to 
the  one  in  which  water  was  electrolyzed  (Fig.  1 8) ,  chlorine 
collects  at  the  anode  and  hydrogen  at  the  cathode.  At 
first  the  chlorine  dissolves  in  the  water,  but  soon  the  water 
in  the  one  tube  becomes  saturated  with  it,  and  if  the  stop- 
cocks are  left  open  until  this  is  the  case,  and  are  then 
closed,  it  will  be  seen  that  the  two  gases  are  set  free  in 
equal  volumes. 


1 84     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

When  measured  volumes  of  the  two  gases  are  caused  to 
unite  it  is  found  that  one  volume  of  hydrogen  combines 
with  one  of  chlorine.  Other  experiments  show  that  the  vol- 
ume of  hydrochloric  acid  formed  is  just  equal  to  the  sum 
of  the  volumes  of  hydrogen  and  chlorine.  Therefore  one 
volume  of  hydrogen  combines  with  one  volume  of  chlorine 
to  form  two  volumes  of  hydrochloric  acid  gas.  Since  chlo- 
rine is  35.18  times  as  heavy  as  hydrogen,  it  follows  that 
one  part  of  hydrogen  by  weight  combines  with  35.18  parts 
of  chlorine  to  form  36.18  parts  of  hydrochloric  acid. 

Physical  properties.  Hydrochloric  acid  is  a  colorless  gas 
which  has  an  irritating  effect  when  inhaled,  and  possesses  a 
sour,  biting  taste,  but  no  marked  odor.  It  is  heavier  than 
air  (density  =  1.26)  and  is  very  soluble  in  water.  Under 
standard  conditions  I  volume  of  water  dissolves  about  500 
volumes  of  the  gas.  On  warming  such  a  solution  the  gas 
escapes,  until  at  the  boiling  point  the  solution  contains 
about  20%  by  weight  of  HC1.  Further  boiling  will  not 
drive  out  any  more  acid,  but  the  solution  will  distill  with 
unchanged  concentration.  A  more  dilute  solution  than  this 
will  lose  water  on  boiling  until  it  has  reached  the  same  con- 
centration, 20%,  and  will  then  distill  unchanged.  Under 
high  pressure  the  gas  can  be  liquefied,  28  atmospheres 
being  required  at  o°.  Under  these  conditions  it  forms  a 
colorless  liquid  which  is  not  very  active  chemically.  It  boils 
at  —80°  and  solidifies  at  —  113°.  The  solution  of  the  gas 
in  water  is  used  almost  entirely  in  the  place  of  the  gas  itself, 
since  it  is  not  only  far  more  convenient  but  also  more  active. 

Chemical  properties.  The  most  important  chemical  prop- 
erties of  hydrochloric  acid  are  the  following : 

i.  Action  as  an  acid.  In  aqueous  solution  hydrochloric 
acid  has  very  strong  acid  properties ;  indeed,  it  is  one  of 


THE  CHLORINE  FAMILY  185 

the  strongest  acids.    It  acts  upon  oxides  and  hydroxides, 
converting  them  into  salts  : 

NaOH  +  HC1  =  NaCl  +  H2O, 
CuO  +  2  HC1  =  CuCl2  +  H2O. 

It  acts  upon  many  metals,  forming  chlorides  and  liberating 
hydrogen:  Zn  +  2  HC1  =  ZnCl2  +  2  H, 

Al  +  3  HC1  =  A1C18  +  3  H. 

Unlike  nitric  and  sulphuric  acids  it  has  no  oxidizing 
action,  so  that  when  it  acts  on  metals  hydrogen  is  always 
given  off. 

2.  Relation  to  combustion.  '  Hydrochloric  acid  gas  is  not 
readily  decomposed,  and  is  therefore  neither  combustible 
nor  a  supporter  of  combustion. 

3.  Action   on  oxidizing  agents.    Although   hydrochloric 
acid  is  incombustible,  it  can  be  oxidized  under  some  cir- 
cumstances, in  which  case  the  hydrogen  combines  with 
oxygen,  while  the  chlorine  is  set  free.    Thus,  when  a  solu- 
tion of  hydrochloric   acid   acts   upon   manganese  dioxide 
part  of  the  chlorine  is  set  free  : 

MnO2  +  4  HC1  =  MnCl2  +  2  H2O  +  2  Cl. 

Aqua  regia.  It  has  been  seen  that  when  nitric  acid  acts 
as  an  oxidizing  agent  it  usually  decomposes,  as  represented 
in  the  equation 

2  HNO3  =  H2O  +  2  NO  4-  3  O. 
The  oxygen  so  set  free  may  act  on  hydrochloric  acid  : 


The  complete  equation  therefore  is 

2  HN03  +  6  HC1  =  4  H20  +  2  NO  +  6  Cl. 


1 86     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

When  concentrated  nitric  and  hydrochloric  acids  are 
mixed  this  reaction  goes  on  slowly,  chlorine  and  some 
other  substances  not  represented  in  the  equation  being 
formed.  The  mixture  is  known  as  aqua  regia  and  is  com- 
monly prepared  by  adding  one  volume  of  nitric  acid  to 
three  volumes  of  hydrochloric  acid.  It  acts  more  powerfully 
upon  metals  and  other  substances  than  either  of  the  acids 
separately,  and  owes  its  strength  not  to  acid  properties 
but.  to  the  action  of  the  nascent  chlorine  which  it  liberates. 
Consequently,  when  it  acts  upon  metals  such  as  gold 'it 
converts  them  into  chlorides,  and  the  reaction  can  be  repre- 
sented by  such  equations  as 

Au  +  3  Cl  =  AuCl8. 

Salts  of  hydrochloric  acid,  —  chlorides.  The  chlorides  of 
all  the  metals  are  known  and  many  of  them  are  very 
important  compounds.  Some  of  them  are  found  in  nature, 
and  all  can  be  prepared  by  the  general  method  of  pre- 
paring salts.  Silver  chloride,  lead  chloride,  and  mercurous 
chloride  are  insoluble  in  water  and  acids,  and  can  be  pre- 
pared by  adding  hydrochloric  acid  to  solutions  of  com- 
pounds of  the  respective  elements.  While  tne  chlorides 
have  formulas  similar  to  the  fluorides,  their  properties  are 
often  quite  different.  This  is  seen  in  the  solubility  of  the 
salts.  Those  metals  whose  chlorides  are  insoluble  form 
soluble  fluorides,  while  many  of  the  metals  which  form 
soluble  chlorides  form  insoluble  fluorides. 

Compounds  of  chlorine  with  oxygen  and  hydrogen.  Chlo- 
rine combines  with  oxygen  and  hydrogen  to  form  four  dif- 
ferent acids.  They  are  all  quite  unstable,  and  most  of  them 
cannot  be  prepared  in  pure  form ;  their  salts  can  easily  be 
made,  however,  and  some  of  them  will  be  met  with  in  the 


THE  CHLORINE  FAMILY 


I87 


study  of  the  metals.    The  formulas  and  names  of  these  acids 
are  as  follows  : 

HC1O hypochlorous  acid. 

HC1O2 chlorous  acid. 


HC10, 


chloric  acid. 


HC1O4 perchloric  acid. 

Oxides  of  chlorine.  Two  oxides  are  known,  having  the  formulas 
C12O  and  C1O2.  They  decompose  very  easily  and  are  good  oxidiz- 
ing agents. 

BROMINE 

Historical.  Bromine  was  discovered  in  1826  by  the 
French  chemist  Ballard,  who  isolated  it  from  sea  salt.  He 
named  it  bromine  (stench)  because  of  its  unbearable  fumes. 

Occurrence.  Bromine  occurs  almost  entirely  in  the  form 
of  bromides,  especially  as  sodium  bromide  and  magnesium 
bromide,  which  are  found  in  many  salt  springs  and  salt 
deposits.  The  Stassfurt 
deposits  in  Germany  and 
the  salt  waters  of  Ohio 
and  Michigan  are  espe- 
cially rich  in  bromides. 

Preparation  of  bro- 
mine. The  laboratory 
method  of  preparing 
bromine  is  essentially 
different  from  the  com- 
mercial method. 

I.  Laboratory  method. 
As  in  the  case  of  chlo-  FlG>55 

rine,  bromine  can  be  prepared  by  the  action  of  hydrobromic 
acid  (HBr)  on  manganese  dioxide.  Since  hydrobromic  acid 
is  not  an  article  of  commerce,  a  mixture  of  sulphuric  acid 


1 88     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

and  a  bromide  is  commonly  substituted  for  it.  The  mate- 
rials are  placed  in  a  retort  arranged  as  shown  in  Fig.  5  5 . 
The  end  of  the  retort  just  touches  the  surface  of  the  water 
in  the  test  tube.  On  heating,  the  bromine  distills  over  and 
is  collected  in  the  cold  receiver.  The  equation  is 

2  NaBr  +  2  H2SO4  +  MnO2  =  Na2SO4  +  MnSO4 
+  2  H2O  +  2  Br. 

2.  Commercial  method.  Bromine  is  prepared  commer- 
cially from  the  waters  of  salt  wells  which  are  especially 
rich  in  bromides.  On  passing  a  current  of  electricity 
through  such  waters  the  bromine  is  first  liberated.  Any 
chlorine  liberated,  however,  will  assist  in  the  reaction, 
since  free  chlorine  decomposes  bromides,  as  shown  in  the 
equation  NaBr+ Cl  =  NaCl  +  Br.  : 

When  the  water  containing  the  bromine  is  heated,  the 
liberated  bromine  distills  over  into  the  receiver. 

Physical  properties.  Bromine  is  a  dark  red  liquid  about 
three  times  as  heavy  as  water.  Its  vapor  has  a  very  offen- 
sive odor  and  is  most  irritating  to  the  eyes  and  throat. 
The  liquid  boils  at  59°  and  solidifies  at  —  7°  ;  but  even  at 
ordinary  temperatures  it  evaporates  rapidly,  forming  a 
reddish-brown  gas  very  similar  to  nitrogen  peroxide  in 
appearance.  Bromine  is  somewhat  soluble  in  water,  100 
volumes  of  water  under  ordinary  conditions  dissolving 
i  volume  of  the  liquid.  It  is  readily  soluble  in  carbon 
disulphide,  forming  a  yellow  solution. 

Chemical  properties  and  uses.  In  chemical  action  bromine 
is  very  similar  to  chlorine.  It  combines  directly  with  many 
of  the  same  elements  with  which  chlorine  unites,  but  with 
less  energy.  It  combines  with  hydrogen  and  takes  away 


THE  CHLORINE  FAMILY 


189 


the  latter  element  from  some  of  its  compounds,  but  not  so 
readily  as  does  chlorine.  Its  bleaching  properties  are  also 
less  marked. 

Bromine  finds  many  uses  in  the  manufacture  of  organic 
drugs  and  dyestuffs  and  in  the  preparation  of  bromides. 

Hydrobromic  acid  (HBr).  When  sulphuric  acid  acts  upon 
a  bromide  hydrobromic  acid  is  set  free : 

2  NaBr  +  H2SO4  =  Na2SO4  +  2  HBr. 

At  the  same  time  some  bromine  is  set  free,  as  may  be  seen 
from  the  red  fumes  which  appear,  and  from  the  odor.  The 
explanation  of  this  is  found  in  the  fact  that  hydrobromic 
acid  is  much  less  stable  than  hydrochloric  acid,  and  is 
therefore  more  easily  oxidized.  Concentrated  sulphuric 
acid  is  a*  good  oxidizing  agent,  and  oxidizes  a  part  of  the 
hydrobromic  acid,  liberating  bromine  : 

H2SO4  +  2  HBr  =  2  H2O  +  SO2  +  2  Br. 

Preparation  of  pure  hydrobromic  acid.  A  convenient 
way  to  make  pure  hydro- 
bromic acid  is  by  the 
action  of  bromine  upon 
moist  red  phosphorus. 
This  can  be  done  with  the 
apparatus  shown  in  Fig.  56. 
Bromine  is  put  into  the 

dropping  funnel    A,    and 
FIG. 56 

red   phosphorus,  together 

with  enough  water  to  cover  it,  is  placed  in  the  flask  B.  By  means  of 
the  stopcock  the  bromine  is  allowed  to  flow  drop  by  drop  into  the 
flask,  the  reaction  taking  place  without  the  application  of  heat.  The 
equations  are 

(1)  P  +  3Br=PBr3, 

(2)  PBr3  +  3  H20  =  P(OH)3  +  3  HBr. 


190     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

The  U-tube  C  contains  glass  beads  which  have  been  moistened 
with  water  and  rubbed  in  red  phosphorus.  Any  bromine  escaping 
action  in  the  flask  acts  upon  the  phosphorus  in  the  U-tube.  The 
hydrobromic  acid  is  collected  in  the  same  way  as  hydrochloric  acid. 

Properties.  Hydrobromic  acid  very  strikingly  resembles 
hydrochloric  acid  in  physical  and  chemical  properties.  It 
is  a  colorless,  strongly  fuming  gas,  heavier  than  hydro- 
chloric acid  and,  like  it,  is  very  soluble  in  water.  Under 
standard  conditions  I  volume  of  water  dissolves  6 10  volumes 
of  the  gas.  Chemically,  the  chief  point  in  which  it  differs 
from  hydrochloric  acid  is  in  the  fact  that  it  is  much  more 
easily  oxidized,  so  that  bromine  is  more  readily  set  free 
from  it  than  chlorine  is  from  hydrochloric  acid. 

Salts  of  hydrobromic  acid,  —  bromides.  The  bromides 
are  very  similar  to  the  chlorides  in  their  properties.  Chlo- 
rine acts  upon  both  bromides  and  free  hydrobromic  acid, 
liberating  bromine  from  them  : 

KBr  +  Cl  =  KC1  +  Br, 
HBr+Cl  =  HCl  +  Br. 

Silver  bromide  is  extensively  used  in  photography,  and 
the  bromides  of  sodium  and  potassium  are  used  as  drugs. 

Oxygen  compounds.  No  oxides  of  bromine  are  surely  known,  and 
bromine  does  not  form  so  many  oxygen  acids  as  chlorine  does.  Salts 
of  hypobromous  acid  (HBrO)  and  bromic  acid  (HBrO3)  are  known. 

IODINE 

Historical.  Iodine  was  discovered  in  1812  by  Courtois 
in  the  ashes  of  certain  sea  plants.  Its  presence  was  re- 
vealed by  its  beautiful  violet  vapor,  and  this  suggested  the 
name  iodine  (from  the  Greek  for  violet  appearance). 

Occurrence.  In  the  combined  state  iodine  occurs  in  very 
small  quantities  in  sea  water,  from  which  it  is  absorbed  by 


THE  CHLORINE  FAMILY 


IQI 


certain  sea  plants,  so  that  it  is  found  in  their  ashes.  It 
occurs  along  with  bromine  in  salt  springs  and  beds,  and 
is  also  found  in  Chili  saltpeter. 

Preparation.  Iodine  may  be  prepared  in  a  number  of 
ways,  the  principal  methods  being  the  following: 

1.  Laboratory  method.    Iodine  can  readily  be  prepared 
in  the  laboratory  from  an  iodide  by  the  method  used  in 
preparing  bromine,  except  that  sodium  iodide  is  substi- 
tuted for  sodium  bromide.    It  can  also  be  made  by  passing 
chlorine  into  a  solution  of  an  iodide. 

2.  Commercial  method.    Commercially  iodine   was    for- 
merly prepared  from  seaweed  (kelp),  but  is  now  obtained 
almost  entirely  from  the  deposits  of  Chili  saltpeter.    The 
crude  saltpeter  is  dissolved  in 

water  ~nd  the  solution  evap- 
orated until  the  saltpeter  crystal- 
lizes. The  remaining  liquors, 
known  as  the  "  mother  liquors," 
contain  sodium  iodate  (NaIO3), 
in  which  form  the  iodine  is 
present  in  the  saltpeter.  The 
chemical  reaction  by  which  the 
iodine  is  liberated  from  this 
compound  is  a  complicated  one, 
depending  on  the  fact  that 
sulphurous  acid  acts  upon  iodic  acid,  setting  iodine  free. 
This  reaction  is  shown  as  follows : 

2  HIO3  +  5H2SO3  -  5  H2SO4  +  H2O  +  2  I. 

Purification  of  iodine.  Iodine  can  be  purified  very  conveniently  in 
the  following  way.  The  crude  iodine  is  placed  in  an  evaporating 
dish  E  (Fig.  57),  and  the  dish  is  set  upon  the  sand  bath. S.  The 
iodine  is  covered  with  the  inverted  funnel  /%  and  the  sand  bath  is 


FIG. 57 


192     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

gently  heated  with  a  Bunsen  burner.  As  the  dish  becomes  warm 
the  iodine  rapidly  evaporates  and  condenses  again  on  the  cold  surface 
of  the  funnel  in  shining  crystals. 

This  process,  in  which  a  solid  is  converted  into  a  vapor  and  is 
again  condensed  into  a  solid  without  passing  through  the  liquid 
state,  is  called  sublimation. 

Physical  properties.  Iodine  is  a  purplish-black,  shining, 
heavy  solid  which  crystallizes  in  brilliant  plates.  Even  at 
ordinary  temperatures  it  gives  off  a  beautiful  violet  vapor, 
which  increases  in  amount  as  heat  is  applied.  It  melts 
at  107°  and  boils  at  175°.  It  is  slightly  soluble  in  water, 
but  readily  dissolves  in  alcohol,  forming  a  brown  solution 
(tincture  of  iodine),  and  in  carbon  disulphide,  forming  a 
violet  solution.  The  element  has  a  strong,  unpleasant  odor, 
though  by  no  means  as  irritating  as  that  of  chlorine  and 
bromine. 

Chemical  properties.  Chemically  iodine  is  quite  similar  to 
chlorine  and  bromine,  but  is  still  less  active  than  bromine. 
It  combines  directly  with  many  elements  at  ordinary  tem- 
peratures. At  elevated  temperatures  it  combines  with 
hydrogen,  but  the  reaction  is  reversible  and  the  compound 
formed  is  quite  easily  decomposed.  Both  chlorine  and 
bromine  displace  it  from  its  salts  : 

KI  +  Br  =  KBr  +  I, 
KI  +  Cl  =  KC1  +  I. 

When  even  minute  traces  of  iodine  are  added  to  thin  starch 
paste  a  very  intense  blue  color  develops,  and  this  reaction 
forms  a  delicate  test  for  iodine.  Iodine  is  extensively  used 
in  medicine,  especially  in  the  form  of  a  tincture.  It  is  also 
largely  used  in  the  preparation  of  dyes  and  organic  drugs, 
lodoform,  a  substance  used  as  an  antiseptic,  has  the  formula 
CHI,. 


THE  CHLORINE  FAMILY  193 

Hydriodic  acid  (HI).  This  acid  cannot  be  prepared  in 
pure  condition  by  the  action  of  sulphuric  acid  upon  an 
iodide,  since  the  hydriodic  acid  set  free  is  oxidized  by  the 
sulphuric  acid  just  as  in  the  case  of  hydrobromic  acid,  but 
to  a  much  greater  extent.  It  can  be  prepared  in  exactly 
the  same  way  as  hydrobromic  acid,  iodine  being  substituted 
for  bromine.  It  can  also  be  prepared  by  passing  hydrosul- 
phuric  acid  into  water  in  which  iodine  is  suspended.  The 

equation  is 

H2S  +  2  I  =  -2.  HI  +  S. 

The  hydriodic  acid  formed  in  this  way  dissolves  in  the 
water. 

Properties  and  uses.  Hydriodic  acid  resembles  the  cor- 
responding acids  of  chlorine  and  bromine  in  physical  prop- 
erties, being  a  strongly  fuming,  colorless  gas,  readily 
soluble  in  water.  Under  standard  conditions  I  volume  of 
water  dissolves  about  460  volumes  of  the  gas.  It  is,  how- 
ever, more  unstable  than  either  hydrochloric  or  hydrobromic 
acids,  and  on  exposure  to  the  air  it  gradually  decomposes  in 
accordance  with  the  equation 

2HI+O=H2O  +  2l. 

Owing  to  the  slight  affinity  between  iodine  and  hydrogen 
the  acid  easily  gives  up  its  hydrogen  and  is  therefore  a 
strong  reducing  agent.  This  is  seen  in  its  action  on  sul- 
phuric acid. 

The  salts  of  hydriodic  acid,  the  iodides,  are,  in  general, 
similar  to  the  chlorides  and  bromides.  Potassium  iodide 
(KI)  is  the  most  familiar  of  the  iodides  and  is  largely  used 
in  medicine. 

Oxygen  compounds.  Iodine  has  a  much  greater  affinity  for  oxygen 
than  has  either  chlorine  or  bromine.  When  heated  with  nitric  acid 


194     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

it  forms  a  stable  oxide  (I2O5).  Salts  of  iodic  acid  (HIO3)  and  periodic 
acid  (HIO4)  are  easily  prepared,  and  the  free  acids  are  much  more 
stable  than  the  corresponding  acids  of  the  other  members  of  this 
family. 

GAY-LUSSAC'S  LAW  OF  VOLUMES 

In  the  discussion  of  the  composition  of  hydrochloric  acid 
it  was  stated  that  one  volume  of  hydrogen  combines  with 
one  volume  of  chlorine  to  form  two  volumes  of  hydrochloric 
acid.  With  bromine  and  iodine  similar  combining  ratios  hold 
good.  These  facts  recall  the  simple  volume  relations  already 
noted  in  the  study  of  the  composition  of  steam  and  ammonia. 
These  relations  may  be  represented  graphically  in  the  fol- 
lowing way  : 

H  I  +  I  ci  I  =  |H  cii  +  IH  ci 


In  the  early  part  of  the  past  century  Gay-Lussac,  a  dis- 
tinguished French  chemist,  studied  the  volume  relations  of 
many  combining  gases,  and  concluded  that  similar  relations 
always  hold.  His  observations  are  summed  up  in  the  follow- 
ing law :  When  two  gases  combine  chemically  there  is 
always  a  simple  ratio  between  their  volumes,  and  between 
the  volume  of  either  one  of  them  and  that  of  the  product, 
provided  it  is  a  gas.  By  a  simple  ratio  is  meant  of  course 
the  ratio  of  small  whole  numbers,  as  1:2,  2:3. 

EXERCISES 

1.  How  do  we  account  for  the  fact  that  liquid  hydrofluoric  acid 
is  not  an  electrolyte  ? 

2.  Why  does  sulphuric  acid  liberate  hydrofluoric  acid  from  its 
salts? 


THE  CHLORINE  FAMILY  195 

3.  In  the  preparation  of  chlorine,  what  advantages  are  there  in 
treating  manganese  dioxide  with  a  mixture  of  sodium  chloride  and 
sulphuric  acid  rather  than  with  hydrochloric  acid  ? 

4.  Why  must  chlorine  water  be  kept  in  the  dark  ? 

5.  What  is  the  derivation  of  the  word  nascent? 

6.  What  substances  studied  are  used  as  bleaching  agents  ?    To 
what  is  the  bleaching  action  due  in  each  case  ? 

7.  What  substances  studied  are  used  as  disinfecting  agents  ? 

8.  What  is  meant  by  the  statement  that  hydrochloric  acid  is  one 
of  the  strongest  acids  ? 

9.  What  is  the  meaning  of  the  phrase  aqua  regia  ? 

10.  C12O  is  the  anhydride  of  what  acid? 

11.  A  solution  of  hydriodic  acid  on  standing  turns  brown.    How 
is  this  accounted  for  ? 

12.  How  can  bromine  vapor   and   nitrogen  peroxide  be  distin- 
guished from  each  other? 

13.  Write  the  equations  for  the  reaction  taking  place  when  hydri- 
odic acid  is  prepared  from  iodine,  phosphorus,  and  water. 

14.  From  their  behavior  toward  sulphuric  acid,  to  what  class  of 
agents  do  hydrobromic  and  hydriodic  acids  belong? 

15.  Give   the   derivation  of  the  names  of  the  elements  of  the 
chlorine  family. 

16.  Write  the  names  and  formulas  for  the  binary  acids  of  the 
group  in  the  order  of  the  stability  of  the  acids. 

17.  What  is  formed  when  a  metal  dissolves  in  each  of  the  follow- 
ing ?  nitric  acid ;  dilute  sulphuric  acid ;  concentrated  sulphuric  acid ; 
hydrochloric  acid  ;  aqua  regia. 

18.  How  could  you  distinguish  between  a  chloride,  a  bromide, 
and  an  iodide  ? 

19.  What  weight  of  sodium  chloride  is  necessary  to  prepare  suf- 
ficient hydrochloric  acid  to  saturate   I  1.  of  water  under  standard 
conditions? 

20.  On  decomposition  100  1.  of  hydrochloric  acid  would  yield  how 
many  liters  of  hydrogen  and  chlorine  respectively,  the  gases  being 
measured  under  the  same  conditions?    Are  your  results  in  accord 
with  the  experimental  facts  ? 


CHAPTER    XVII 
CARBON  AND  SOME  OF  ITS  SIMPLER  COMPOUNDS 

The  family.  Carbon  stands  at  the  head  of  a  family  of 
elements  in  the  fourth  group  in  the  periodic  table.  The 
resemblances  between  the  elements  of  this  family,  while 
quite  marked,  are  not  so  striking  as  in  the  case  of  the  elements 
of  the  chlorine  family.  With  the  exception  of  carbon,  these 
elements  are  comparatively  rare,  and  need  not  be  taken  up 
in  detail  in  this  chapter.  Titanium  will  be  referred  to  again 
in  connection  with  silicon  which  it  very  closely  resembles. 

Occurrence.  Carbon  is  found  in  nature  in  the  uncom- 
bined  state  in  several  forms.  The  diamond  is  practically 
pure  carbon,  while  graphite  and  coal  are  largely  carbon,  but 
contain  small  amounts  of  other  substances.  Its  natural 
compounds  are  exceedingly  numerous  and  occur  as  gases, 
liquids,  and  solids.  Carbon  dioxide  is  its  most  familiar 
gaseous  compound.  Natural  gas  and  petroleum  are  largely 
compounds  of  carbon  with  hydrogen.  The  carbonates, 
especially  calcium  carbonate,  constitute  great  strata  of 
rocks,  and  are  found  in  almost  every  locality.  All  living 
organisms,  both  plant  and  animal,  contain  a  large  per- 
centage of  this  element,  and  the  number  of  its  compounds 
which  go  to  make  up  all  the  vast  variety  of  animate  nature 
is  almost  limitless.  Over  one  hundred  thousand  definite 
compounds  containing  carbon  have  been  prepared.  In  the 
free  state  carbon  occurs  in  three  allotropic  forms,  two  of 
which  are  crystalline  and  one  amorphous. 

196 


CARBON  AND  SOME  OF  ITS  COMPOUNDS      197 

Crystalline  carbon.  Crystalline  carbon  occurs  in  two 
forms,  —  diamond  and  graphite. 

i.  Diamond.  Diamonds  are  found  in  considerable  quan- 
tities in  several  localities,  especially  in  South  Africa,  the 
East  Indies,  and  Brazil.  The  crystals  belong  to  the  regular 
system,  but  the  natural  stones  do  not  show  this  very  clearly. 
When  found  they  are  usually  covered  with  a  rough  coating 
which  is  removed  in  the  process  of  cutting.  Diamond  cut- 
ting is  carried  on  most  extensively  in  Holland. 

The  density  of  the  diamond  is  3.5,  and,  though  brittle, 
it  is  one  of  the  hardest  of  substances.  Black  diamonds, 
as  well  as  broken  and  imperfect  stones  which  are  valueless 
as  gems,  are  used  for  grinding  hard  substances.  Few  chem- 
ical reagents  have  any  action  on  the  diamond,  but  when 
heated  in  oxygen  or  the  air  it  blackens  and  burns,  forming 
carbon  dioxide. 

Lavoisier  first  showed  that  carbon  dioxide  is  formed  by 
the  combustion  of  the  diamond ;  and  Sir  Humphry  Davy 
in  1814  showed  that  this  is  the  only  product  of  combustion, 
and  that  the  diamond  is  pure  carbon. 

The  diamond  as  a  gem.  The  pure  diamond  is  perfectly  transpar- 
ent and  colorless,  but  many  are  tinted  a  variety  of  colors  by  traces 
of  foreign  substances.  Usually  the  colorless  ones  are  the  most  highly 
prized,  although  in  some  instances  the  color  adds  to  the  value  ;  thus 
the  famous  Hope  diamond  is  a  beautiful  blue.  Ijght  passing  through 
a  diamond  is  very  much  refracted,  and  to  this  fact  the  stone  owes 
its  brilliancy  and  sparkle. 

Artificial  preparation  of  diamonds.  Many  attempts  have  been 
made  to  produce  diamonds  artificially,  but  for  a  long  time  these 
always  ended  in  failure,  graphite  and  not  diamonds  being  the  product 
obtained.  The  French  chemist  Moissan,  in  his  extended  study  of 
chemistry  at  high  temperatures,  finally  succeeded  (1893)  in  making 
some  small  ones.  He  accomplished  this  by  dissolving  carbon  in  boil- 
ing iron  and  plunging  the  crucible  containing  the  mixture  into  water, 


198     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

as  shown  in  Fig.  58.    Under  these  conditions  the  carbon  crystallized 
in  the  iron  in  the  form  of  the  diamond.    The  diamonds  were  then 

obtained  by  dissolving  away  the  iron 

in  hydrochloric  acid. 

2.  Graphite.  This  form  of  car- 
bon is  found  in  large  quantities, 
especially  in  Ceylon,  Siberia,  and 
in  some  localities  of  the  United 
States  .and  Canada.  It  is  a 
shining  black  substance,  very 
soft  and  greasy  to  the  touch. 
8  Its  density  is  about  2.15.  It 

varies    somewhat    in    properties 

according  to  the  locality  in  which  it  is  found,  and  is  more 
easily  attacked  by  reagents  than  is  the  diamond.  It  is  also 
manufactured  by  heating  carbon  with  a  small  amount  of 
iron  (3%)  in  an  electric  furnace.  It  is  used  in  the  manu- 
facture of  lead  pencils  and  crucibles,  as  a  lubricant,  and  as  a 
protective  covering  for  iron  in  the  form  of  a  polish  or  a  paint. 
Amorphous  carbon.  Although  there  are  many  varieties 
of  amorphous  carbon  known,  they  are  not  true  allotropic- 
modifications.  They  differ  merely  in  their  degree  of  purity, 
their  fineness  of  division,  and  in  their  mode  of  preparation. 
These  substances  are  of  the  greatest  importance,  owing  to 
their  many  uses  in  the  arts  and  industries.  As  they  occur 
in  nature,  or  are  made  artificially,  they  are  nearly  all  impure 
carbon,  the  impurity  depending  on  the  particular  substance 
in  question. 

i.  Pure  carbon.  Pure  amorphous  carbon  is  best  pre- 
pared by  charring  sugar.  This  is  a  substance  consisting  of 
carbon,  hydrogen,  and  oxygen,  the  latter  two  elements  being 
present  in  the  ratio  of  one  oxygen  atom  to  two  of  hydrogen. 


CARBON  AND  SOME  OF  ITS  COMPOUNDS      199 

When  sugar  is  strongly  heated  the  oxygen  and  hydrogen 
are  driven  off  in  the  form  of  water  and  pure  carbon  is  left 
behind.  Prepared  in  this  way  it*  is  a  soft,  lustrous,  very 
bulky,  black  powder. 

2.  Coal  and  coke.    Coals  of  various  kinds  were  probably 
formed  from  vast  accumulations  of  vegetable  matter  in 
former  ages,  which  became  covered  over  with  earthy  mate- 
rial and  were   thus  protected  from   rapid   decay.    Under 
various   natural   agencies   the  organic  matter  was  slowly 
changed  into  coal.    In  anthracite  these  changes  have  gone 
the  farthest,  and  this  variety  of  coal  is  nearly  pure  carbon. 
Soft  or  bituminous  coals  contain  considerable  organic  mat- 
ter besides  carbon  and  mineral  substances.    When  heated 
strongly  out    of   contact  with   air  the   organic  matter  is 
decomposed  and  the  resulting  volatile  matter  is  driven  off 
in  the  form  of  gases  and   vapors,  and    only  the  mineral 
matter  and  carbon  remain  behind.    The  gaseous  product 
is  chiefly  illuminating  gas  and  the  solid  residue  is  coke. 
Some  of   the  coke  is  found  as  a  dense  cake  on  the  sides 
and  roof  of  the  retort.     This  is  called  retort  carbon  and 
is  quite  pure. 

3.  Charcoal.    This  is  prepared  from  wood  in  the  same 
way  that  coke  is  made  from  coal.    When  the  process  is 
carried  on  in  retorts  the  products  expelled  by  the  heat  are 
saved.    Among  these  are  many  valuable*  substances  such 
as  wood  alcohol  and  acetic  acid.    Where  timber  is  abun- 
dant the  process  is  carried  out  in  a  wasteful  way,  by  merely 
covering  piles  of  wood  with  sod  and  setting  the  wood  on 
fire.    Some  wood  burns  and  the  heat  from  this  decomposes 
the  wood  not  burned,  forming  charcoal  from  it.    The  char- 
coal, of  course,  contains  the  mineral  part  of  the  wood  from 
which  it  is  formed. 


200     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

4.  Bone  black.    This  is  sometimes  called  animal  charcoal, 
and  is  made  by  charring  bones  and  animal  refuse.    The 
organic  part  of  the  materials  is  thus  decomposed  and  car- 
bon is  left  in  a  very  finely  divided  state,  scattered  through 
the  mineral  part  which  consists  largely  of  calcium  phos- 
phate.   For  some  uses  this  mineral  part  is  removed  by 
treatment  with  hydrochloric  acid  and  prolonged  washing. 

5.  Lampblack.    Lampblack    and    soot    are    products    of 
imperfect  combustion  of  oil  and  coal,  and  are  deposited 
from  a  smoky  flame  on  a  cold  surface.    The  carbon  in  this 
form  is  very  finely  divided  and   usually  contains  various 
oily  materials. 

Properties.  While  the  various  forms  of  carbon  differ  in 
many  properties,  especially  in  color  and  hardness,  yet  they 
are  all  odorless,  tasteless  solids,  insoluble  in  water  and 
characterized  by  their  stability  towards  heat.  Only  in  the 
intense  heat  of  the  electric  arc  does  carbon  volatilize,  pass- 
ing directly  from  the  solid  state  into  a  vapor.  Owing  to 
this  fact  the  inside  surface  of  an  incandescent  light  bulb 
after  being  used  for  some  time  becomes  coated  with  a  dark 
film  of  carbon.  It  is  not  acted  on  at  ordinary  temperatures 
by  most  reagents,  but  at  a  higher  temperature  it  combines 
directly  with  many  of  the  elements,  forming  compounds 
called  carbides.  When  heated  in  the  presence  of  sufficient 
oxygen  it  burns,  forming  carbon  dioxide. 

Uses  of  carbon.  The  chief  use  of  amorphous  carbon  is 
for  fuel  to  furnish  heat  and  power  for  all  the  uses  of  civi- 
lization. An  enormous  quantity  of  carbon  in  the  form  of 
the  purer  coals,  coke,  and  charcoal  is  used  as  a  reducing 
agent  in  the  manufacture  of  the  various  metals,  especially 
in  the  metallurgy  of  iron.  Most  of  the  metals  are  found 
in  nature  as  oxides,  or  in  forms  which  can  readily  be 


CARBON  AND  SOME  OF  ITS  COMPOUNDS      2OI 

converted  into  oxides.  When  these  oxides  are  heated  with 
carbon  the  oxygen  is  abstracted,  leaving  the  metal.  Retort 
carbon  and  coke  are  used  to  make  electric  light  carbons 
and  battery  plates,  while  lampblack  is  used  for  indelible 
inks,  printer's  ink,  and  black  varnishes.  Bone  black  and 
charcoal  have  the  property  of  absorbing  large  volumes  of 
certain  gases,  as  well  as  smaller  amounts  of  organic  matter; 
hence  they  are  used  in  filters  to  remove  noxious  gases  and 
objectionable  colors  and  odors  from  water.  Bone  black  is 
used  extensively  in  the  sugar  refineries  to  remove  coloring 
matter  from  the  impure  sugars. 

Chemistry  of  carbon  compounds.  Carbon  is  remarkable 
for  the  very  large  number  of  compounds  which  it  forms 
with  the  other  elements,  especially  with  oxygen  and  hydro- 
gen. Compounds  containing  carbon  are  more  numerous 
than  all  others  put  together,  and  the  chemistry  of  these 
substances  presents  peculiarities  not  met  with  in  the  study 
of  other  substances.  For  these  reasons  the  systematic 
study  of  carbon  compounds,  or  of  organic  chemistry  as  it 
is  usually  called,  must  be  deferred  until  the  student  has 
gained  some  knowledge  of  the  chemistry  of  other  elements. 
An  acquaintance  with  a  few  of  the  most  familiar  carbon 
compounds  is,  however,  essential  for  the  understanding  of 
the  general  principles  of  chemistry. 

Compounds  of  carbon  with  hydrogen,  — *he  hydrocarbons. 
Carbon  unites  with  hydrogen  to  form  a  very  large  number 
of  compounds  called  hydrocarbons.  Petroleum  and  natural 
gas  are  essentially  mixtures  of  a  great  variety  of  these 
hydrocarbons.  Many  others  are  found  in  living  plants, 
and  still  others  are  produced  by  the  decay  of  organic 
matter  in  the  absence  of  air.  Only  two  of  them,  methane 
and  acetylene,  will  be  discussed  here. 


202     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Methane  (marsh  gas]  (CH4).  This  is  one  of  the  most 
important  of  these  hydrocarbons,  and  constitutes  about 
nine  tenths  of  natural  gas.  As  its  name  suggests,  it  is 
formed  in  marshes  by  the  decay  of  vegetable  matter  under 
water,  and  bubbles  of  the  gas  are  often  seen  to  rise  when 
the  dead  leaves  on  the  bottom  of  pools  are  stirred.  It  also 
collects  in  mines,  and,  when  mixed  with  air,  is  called  fire 
damp  by  the  miners  because  of  its  great  inflammability, 
damp  being  an  old  name  for  a  gas.  It  is  formed  when 
organic  matter,  such  as  coal  or  wood,  is  heated  in  closed 
vessels,  and  is  therefore  a  principal  constituent  of  coal 
gas. 

Preparation.  Methane  is  prepared  in  the  laboratory  by  heating 
sodium  or  calcium  acetate  with  soda-lime.  Equal  weights  of  fused 
sodium  acetate  and  soda-lime  are  thoroughly  dried,  then  mixed  and 
placed  in  a  good-sized,  hard-glass  test  tube  fitted  with  a  one-holed 
stopper  and  delivery  tube.  The  mixture  is  gradually  heated,  and 
when  the  air  has  been  displaced  from  the  tube  the  gas  is  collected 
in  bottles  by  displacement  of  water.  Soda-lime  is  a  mixture  of 
sodium  and  calcium  hydroxides.  Regarding  it  as  sodium  hydroxide 
alone,  the  equation  is 

NaC2H3O2  +  NaOH  =  Na2CO3  +  CH4. 

Properties.  Methane  is  a  colorless,  odorless  gas  whose 
density  is  0.55.  It  is  difficult  to  liquefy,  boiling  at  —  155° 
under  standard  pressure,  and  is  almost  insoluble  in  water. 
It  burns  with  a  pale  blue  flame,  liberating  much  heat,  and 
when  mixed  with  oxygen  is  very  explosive. 

Davy's  safety  lamp.  In  1815  Sir  Humphry  Davy  invented  a 
lamp  for  the  use  of  miners,  to  prevent  the  dreadful  mine  explosions 
then  common,  due  to  methane  mixed  with  air.  The  invention  con- 
sisted in  surrounding  the  upper  part  of  the  common  miner's  lamp 
with  a  mantle  of  wire  gauze  and  the  lower  part  with  glass  (Fig.  59). 
It  has  been  seen  that  two  gases  will  not  combine  until  raised  to  their 


CARBON  AND  SOME  OF  ITS  COMPOUNDS      203 


kindling  temperature,  and  if  while  combining  they  are  cooled  below 
this  point,  the  combination  ceases.  A  flame  will  not  pass  through 
a  wire  gauze  because  the  metal,  being  a  good  con- 
ductor of  heat,  takes  away  so  much  heat  from  the 
flame  that  the  gases  are  cooled  below  the  kindling 
temperature.  When  a  lamp  so  protected  is  brought 
into  an  explosive  mixture  the  gases  inside  the  wire 
mantle  burn  in  a  series  of  little  explosions,  giving 
warning  to  the  miner  that  the  air  is  unsafe. 

Acetylene  (C2H2).  This  is  a  colorless  gas 
usually  having  a  disagreeable  odor  due  to  im- 
purities. It  is  now  made  in  large  quantities 
from  calcium  carbide  (CaC2).  This  substance 
is  formed  when  coal  and  lime  are  heated  to- 
gether in  an  electric  furnace.  When  treated 


FIG.  59 


with  water  the  carbide  is  decomposed,  yielding  acetylene : 
CaC2  +  2  H20  =  C2H2  +  Ca(OH)2. 

Under  ordinary  conditions  the  gas  burns  with  a  very  smoky 
flame ;  in  burners  constructed  so  as  to  secure  a 
large  amount  of  oxygen  it  burns  with  a  very 
brilliant  white  light,  and  hence  is  used  as  an 
illuminant. 

Laboratory  preparation.  The  gas  can  be  prepared 
readily  in  a  generator  such  as  is  shown  in  Fig.  60. 
The  inner  tube  contains  fragments  of  calcium  carbide, 
while  the  outer  one  is  filled  with*  water.  As  long  as 
the  stopcock  is  closed  the  water  cannot  rise  in  the 
inner  tube.  When  the  stopcock  is  open  the  water 
rises,  and,  coming  into  contact  with  the  carbide  in  the 
inner  tube,  generates  acetylene.  This  escapes  through 
the  stopcock,  and  after  the  air  has  been  expelled  may 
be  lighted  as  it  issues  from  the  burner. 

Carbon  forms  two  oxides,  namely,  carbon  dioxide  (CO2) 
and  carbon  monoxide  (CO). 


FIG. 60 


204     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Carbon  dioxide  (CO2).  Carbon  dioxide  is  present  in  the 
air  to  the  extent  of  about  3  parts  in  10,000,  and  this  appar- 
ently small  amount  is  of  fundamental  importance  in  nature. 
In  some  localities  it  escapes  from  the  earth  in  great  quan- 
tities, and  many  spring  waters  carry  large  amounts  of  it 
in  solution.  When  these  highly  charged  spring  waters 
reach  the  surface  of  the  earth,  and  the  pressure  on  them 
is  removed,  the  carbon  dioxide  escapes  with  effervescence. 
It  is  a  product  of  the  oxidation  of  all  organic  matter,  and 
is  therefore  formed  in  fires  as  well  as  in  the  process  of 
decay.  It  is  thrown  off  from  the  lungs  of  all  animals  in 
respiration,  and  is  a  product  of  many  fermentation  processes 
such  as  vinegar  making  and  brewing.  Combined  with  me- 
tallic oxides  it  forms  vast  deposits  of  carbonates  in  nature. 

Preparation.  In  the  laboratory  carbon  dioxide  is  always 
prepared  by  the  action  of  an  acid  upon  a  carbonate,  usually 
calcium  carbonate,  the  apparatus  shown  in  Fig.  39  serving 
the  purpose  very  well.  This  reaction  might  be  expected 
to  produce  carbonic  acid,  thus  : 

CaCO3  +  2  HC1  =  CaCl2  +  H2CO3. 

Carbonic  acid  is  very  unstable,  however,  and  decomposes 
into  its  anhydride,  CO2,  and  water,  thus  : 

H2C03  =  H20  +  C02. 

The  complete  reaction  is  represented  by  the  equation 
CaCO3  +  2  HC1  =  CaCl2  +  CO2  +  H2O. 

Physical  properties.  Carbon  dioxide  is  a  colorless,  prac- 
tically odorless  gas  whose  density  is  1.5.  Its  weight  may  be 
inferred  from  the  fact  that  it  can  be  siphoned,  or  poured 
like  water,  from  one  vessel  downward  into  another.  At  1 5° 


CARBON  AND  SOME  OF  ITS  COMPOUNDS      205 

and  under  ordinary  pressure  it  dissolves  in  its  own  volume 
of  water  and  imparts  a  somewhat  biting,  pungent  taste  to 
it.  It  is  easily  condensed,  and  is  now  prepared  commercially 
in  this  form  by  pumping  the  gas  into  steel  cylinders  (see 
Fig.  6)  which  are  kept  cold  during  the  process.  When  the 
liquid  is  permitted  to  escape  into  the  air  part  of  it  instantly 
evaporates,  and  in  so  doing  absorbs  so  much  heat  that 
another  portion  is  solidified,  the  solid  form  strikingly 
resembling  snow  in  appearance.  This  snow  is  very  cold 
and  mercury  can  easily  be  frozen  with  it. 

Solid  carbon  dioxide.  Cylinders  of  liquid  carbon  dioxide  are  inex- 
pensive, and  should  be  available  in  every  school.  To  demonstrate 
the  properties  of  solid  carbon  dioxide,  the  cylinder  should  be  placed 
across  the  table  and  supported  in  such  a  way  that  the  stopcock  end 
is  several  inches  lower  than  the  other  end.  A  loose  bag  is  made  by 
holding  the  corners  of  a  handkerchief  around  the  neck  of  the  stop- 
cock, and  the  cock  is  then  turned  on  so  that  the  gas  rushes  out  in 
large  quantities.  Very  quickly  a  considerable  quantity  of  the  snow 
collects  in  the  handkerchief.  To  freeze  mercury,  press  a  piece  of 
filter  paper  into  a  small  evaporating  dish  and  pour  the  mercury 
upon  it.  Coil  a  flat  spiral  upon  the  end  of  a  wire,  and  dip  the  spiral 
into  the  mercury.  Place  a  quantity  of  solid  carbon  dioxide  upon  the 
mercury  and  pour  10  cc.-i5  cc.  of  ether  over  it.  In  a  minute  or  two 
the  mercury  will  solidify  and  may  be  removed  from  the  dish  by 
the  wire  serving  as  a  handle.  The  filter  paper  is  to  prevent  the  mer- 
cury from  sticking  to  the  dish  ;  the  ether  dissolves  the  solid  carbon 
dioxide  and  promotes  its  rapid  conversion  into  gas. 

Chemical  properties.  Carbon  dioxide  is  incombustible, 
since  it  is,  like  water,  a  product  of  combustion.  It  does  not 
support  combustion,  as  does  nitrogen  peroxide,  because  the 
oxygen  in  it  is  held  in  very  firm  chemical  union  with  the 
carbon.  Very  strong  reducing  agents,  such  as  highly  heated 
carbon,  can  take  away  half  of  its  oxygen  : 


206     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Uses.  The  relation  of  carbon  dioxide  to  plant  life 
has  been  discussed  in  a  previous  chapter.  Water  highly 
charged  with  carbon  dioxide  is  used  for  making  soda  water 
and  similar  beverages.  Since  it  is  a  non-supporter  of  com- 
bustion and  can  be  generated  readily,  carbon  dioxide  is 
also  used  as  a  fire  extinguisher.  Some  of  the  portable 
fire  extinguishers  are  simply  devices  for  generating  large 
amounts  of  the  gas.  It  is  not  necessary  that  all  the  oxy- 
gen should  be  kept  away  from  the  fire  in  order  to  smother 
it.  A  burning  candle  is  extinguished  in  air  which  contains 
only  2.5%  of  carbon  dioxide. 

Carbonic  acid  (H2CO3).  Like  most  of  the  oxides  of  the 
non-metallic  elements,  carbon  dioxide  is  an  acid  anhydride. 
It  combines  with  water  to  form  an  acid  of  the  formula 
H2CO3,  called  carbonic  acid  : 

H2O  +  CO2  =  H2CO3. 

The  acid  is,  however,  very  unstable  and  cannot  be  isolated. 
Only  a  very  small  amount  of  it  is  actually  formed  when 
carbon  dioxide  is  passed  into  water,  as  is  evident  from  the 
small  solubility  of  the  gas.  If,  however,  a  base  is  present 
in  the  water,  salts  of  carbonic  acid  are  formed,  and  these 
are  quite  stable  : 

2  NaOH  +  H2O  +  CO2  =  Na2CO3  +  2  H2O. 

Action  of  carbon  dioxide  on  bases.  This  conduct  is  explained  by 
the  principles  of  reversible  reactions.  The  equation 


is  a  reversible  equation,  and  the  extent  to  which  the  reaction  pro- 
gresses depends  upon  the  relative  concentrations  of  each  of  the 
three  factors  in  it.  Equilibrium  is  ordinarily  reached  when  very 
little  H2CO8  is  formed.  If  a  base  is  present  in  the  water  to  com- 
bine with  the  H2CO8  as  fast  as  it  is  formed,  all  of  the  CO2  is  con- 
verted into  HoCOg,  and  thence  into  a  carbonate. 


CARBON  AND  SOME  OF  ITS  COMPOUNDS      207 


Salts  of  carbonic  acid,  —  carbonates.  The  carbonates 
form  a  very  important  class  of  salts.  They  are  found  in 
large  quantities  in  nature,  and  are  often  used  in  chemical 
processes.  Only  the  carbonates  of  sodium,  potassium, 
and  ammonium  are  soluble,  and  these  can  be  made  by 
the  action  of  carbon  dioxide  on  solutions  of  the  bases, 
as  has  just  been  explained. 

The  insoluble  carbonates  are  formed  as  precipitates 
when  soluble  salts  are  treated  with  a  solution  of  a  soluble 
carbonate.  Thus  the  insoluble  calcium  carbonate  can  be 
made  by  bringing  together  solutions  of  calcium  chloride 
and  sodium  carbonate: 

CaCl2  +  Na2CO3  -  CaCO3  +  2  NaCl. 

Most  of  the  carbonates  are  decomposed  by  heat,  yielding  an 
oxide  of  the  metal  and  carbon  dioxide.  Thus  lime  (calcium 
oxide)  is  made  by  strongly  heating  calcium  carbonate: 

CaCO3  =  CaO  4-  CO2. 

Acid  carbonates.  Like  all  acids  containing  two  acid 
hydrogen  atoms,  carbonic  acid  can  form  both  normal  and 
acid  salts.  The  acid  carbonates  are  made  by  treating  a 
normal  carbonate  with  an  excess  of  carbonic  acid.  With 
few  exceptions  they  are  very  unstable,  heat  decomposing 
them  even  when  in  solution. 

Action  of  carbon  dioxide  on  calcium  hydroxide.  If  carbon  dioxide 
is  passed  into  clear  lime  water,  calcium  carbonate  is  at  first  pre- 

cipitated : 

H20  +  CO2  =  H2CO3, 

Ca(OH)2  +  H2CO3  =  CaCO3  +  2  H2O. 

Advantage  is  taken  of  this  reaction  in  testing  for  the  presence  of 
carbon  dioxide,  as  already  explained  in  the  chapter  on  the  atmos- 
phere. If  the  current  of  carbon  dioxide  is  continued,  the  precipitate 


208     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

soon  dissolves,  because  the  excess  of  carbonic  acid  forms  calcium 
acid  carbonate  which  is  soluble  : 

CaCO3  +  H2CO3  =  Ca(HC03)2. 

If  now  the  solution  is  heated,  the  acid  carbonate  is  decomposed  and 
calcium  carbonate  once  more  precipitated  : 

Ca(HCO3)2  =  CaC03  +  H2CO3. 

Carbon  monoxide  (CO).  Carbon  monoxide  can  be  made 
in  a  number  of  ways,  the  most  important  of  which  are 
the  three  following  : 

i.  By  the  partial  oxidation  of  carbon.  If  a  slow  current 
of  air  is  conducted  over  highly  heated  carbon,  the  monox- 
ide is  formed,  thus  :  /-  ,  Q  _ 


It  is  therefore  often  formed  in  stoves  when  the  air  draught 
is  insufficient.  Water  gas,  which  contains  large  amounts 
of  carbon  monoxide,  is  made  by  partially  oxidizing  carbon 
with  steam:  C  +  H2O  =  CO  +  2  H. 

2.  By  the  partial  reduction    of  carbon    dioxide.    When 
carbon  dioxide  is  conducted  over  highly  heated  carbon  it 
is  reduced  to  carbon  monoxide  by  the  excess  of  carbon  : 

CO2  +  C  =  2  CO. 

When  coal  is  burning  in  a  stove  or  grate  carbon  dioxide  is 
at  first  formed  in  the  free  supply  of  air,  but  as  the  hot  gas 
rises  through  the  glowing  coal  it  is  reduced  to  carbon 
monoxide.  When  the  carbon  monoxide  reaches  the  free 
air  above  the  coal  it  takes  up  oxygen  to  form  carbon  diox- 
ide, burning  with  the  blue  flame  so  familiar  above  a  bed 
of  coals,  especially  in  the  case  of  hard  coals. 

3.  By  the  decomposition  of  oxalic  acid.    In  the  laboratory 
carbon    monoxide    is    usually  prepared   by  the   action  of 


CARBON  AND  SOME  OF  ITS  COMPOUNDS      209 


concentrated  sulphuric  acid  upon  oxalic  acid.  The  latter 
substance  has  the  formula  C2H2O4.  The  sulphuric  acid, 
owing  to  its  affinity  for  water,  decomposes  the  oxalic  acid, 
as  represented  in  the  equation 

C2H204  +  (H2S04)  =  (H2S04)  +  H20  +  CO2  +  CO. 

Properties.  Carbon  monoxide  is  a  light,  colorless,  almost 
odorless  gas,  very  difficult  to  liquefy.  Chemically  it  is 
very  active,  combining  directly  with  a  great  many  sub- 
stances. It  has  a  great  affinity  for  oxygen  and  is  therefore 
combustible  and  a  good  reducing  agent.  Thus,  if  carbon 
monoxide  is  passed  over  hot  copper  oxide,  the  copper  is 
reduced  to  the  metallic  state  : 

CuO  +  CO  =  Cu  +  CO2. 

When  inhaled  it  combines  with  the  red  coloring  matter  of 
the  blood  and  in  this  way  prevents  the  absorption  of 
oxygen,  so  that  even  a  small  quantity 
of  the  gas  may  prove  fatal. 

The   reducing   power  of   carbon   monoxide. 
Fig.    6 1     illustrates    a    method   of   showing 

E 


FIG. 6i 

the  reducing  power  of  carbon  monoxide.  The  gas  is  generated  by 
gently  heating  7  or  8  g.  of  oxalic  acid  with  25  cc.  of  concentrated  sul- 
phuric acid  in  a  200  cc.  flask  A.  The  bottle  B  contains  a  solution 
of  sodium  hydroxide,  which  removes  the  carbon  dioxide  formed  along 
with  the  monoxide.  C  contains  a  solution  of  calcium  hydroxide  to 
show  that  the  carbon  dioxide  is  completely  removed.  E  is  a  hard- 
glass  tube  containing  i  or  2  g.  of  copper  oxide,  which  is  heated  by  a 


210     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

burner.  The  black  copper  oxide  is  reduced  to  reddish  metallic  cop- 
per by  the  carbon  monoxide,  which  is  thereby  changed  to  carbon 
dioxide.  The  presence  of  the  carbon  dioxide  is  shown  by  the  pre- 
cipitate in  the  calcium  hydroxide  solution  in  D.  Any  unchanged 
carbon  monoxide  is  collected  over  water  in  F. 

Carbon  disulphide  (CS2).  Just  as  carbon  combines  with 
oxygen  to  form  carbon  dioxide,  so  it  combines  with  sulphur 
to  form  carbon  disulphide  (CS2).  This  compound  has  been 
described  in  the  chapter  on  sulphur. 

Hydrocyanic  acid  (prussic  acid}  (H CN) .  Under  the  proper 
conditions  carbon  unites  with  nitrogen  and  hydrogen  to 
form  the  acid  HCN,  called  hydrocyanic  acid.  It  is  a  weak, 
volatile  acid,  and  is  therefore  easily  prepared  by  treating 
its  salts  with  sulphuric  acid  : 

KCN  +  H2SO4  =  KHSO4  +  HCN. 

It  is  most  familiar  as  a  gas,  though  it  condenses  to  a  color- 
less liquid  boiling  at  26°.  It  has  a  peculiar  odor,  suggesting 
bitter  almonds,  and  is  extremely  poisonous  either  when 
inhaled  or  when  taken  into  the  stomach.  A  single  drop 
may  cause  death.  It  dissolves  readily  in  water,  its  solution 
being  commonly  called  prussic  acid. 

The  salts  of  hydrocyanic  acid  are  called  cyanides,  the 
cyanides  of  sodium  and  potassium  being  the  best  known. 
These  are  white  solids  and  are  extremely  poisonous. 

Solutions  of  potassium  cyanide  are  alkaline.  A  solution 
of  potassium  cyanide  turns  red  litmus  blue,  and  must  there- 
fore contain  hydroxyl  ions.  The  presence  of  these  ions  is 
accounted  for  in  the  following  way. 

Although  water  is  so  little  dissociated  into  its  ions  H+ 
and  OH~  that  for  most  purposes  we  may  neglect  the  dis- 
sociation, it  is  nevertheless  measurably  dissociated.  Hydro- 
cyanic acid  is  one  of  the  weakest  of  acids,  and  dissociates 


CARBON  AND  SOME  OF  ITS  COMPOUNDS      2  1  1 

to  an  extremely  slight  extent.  When  a  cyanide  such  as 
potassium  cyanide  dissolves  it  freely  dissociates,  and  the 
CN~  ions  must  come  to  an  equilibrium  with  the  H+  ions 
derived  from  the  water  : 


The  result  of  this  equilibrium  is  that  quite  a  number  of 
H+  ions  from  the  water  are  converted  into  undissociated 
HCN  molecules.  But  for  every  H+  ion  so  removed  an 
OH~  ion  remains  free,  and  this  will  give  the  solution  alka- 
line properties. 

EXERCISES 

1.  How  can  you  prove  that  the  composition  of  the  different  allo- 
tropic  forms  of  carbon  is  the  same  ? 

2.  Are  lampblack  and  bone  black  allotropic  forms  of  carbon? 
Will  equal  amounts  of  heat  be  liberated  in  the  combustion  of  I  g. 
of  each  ? 

3.  How  could  you  judge  of  the  relative  purity  of  different  forms 
of  carbon  ? 

4.  Apart  from  its  color,  why  should  carbon  be  useful  in  the  prep- 
aration of  inks  and  paints  ? 

5.  Could  asbestos  fibers  be  used  to  replace  the  wire  in  a  safety 
lamp? 

6.  Why  do  most  acids  decompose  carbonates? 

7.  What  effect  would  doubling  the  pressure  have  upon  the  solu- 
bility of  carbon  dioxide  in  water  ? 

8.  What  compound  would  be  formed  by  passing  carbon  dioxide 
into  a  solution  of  ammonium  hydroxide  ?    Write  the  equation. 

9.  Write  equations  for  the  preparation  of  K2CO3  ;  of  BaCO3  ; 
of  MgC03. 

10.  In  what  respects  are  carbonic  and  sulphurous  acids  similar? 

11.  Give  three  reasons  why  the  reaction  which  takes  place  when 
a  solution  of  calcium  acid  carbonate  is  heated,  completes  itself. 

12.  How  could  you  distinguish  between  carbonates  and  sulphites  ? 

13.  How  could  you  distinguish  between  oxygen,  hydrogen,  nitro- 
gen, nitrous  oxide,  and  carbon  dioxide  ? 


212     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

14.  Could  a  solution  of  sodium  hydroxide  be  substituted  for  the 
solution  of  calcium  hydroxide  in  testing  for  carbon  dioxide  ? 

15.  What  weight  of  sodium  hydroxide  is  necessary  to  neutralize 
the  carbonic  acid  formed  by  the  action  of  hydrochloric  acid  on  100  g. 
of  calcium  carbonate  ? 

16.  What  weight  of  calcium  carbonate  would  be  necessary  to 
prepare  sufficient  carbon  dioxide  to  saturate  10  1.  of  water  at  15°  and 
under  ordinary  pressure  ? 

17.  On  the  supposition  that  calcium  carbide  costs  12  cents  a" 
kilogram, -what  would  be  the  cost  of  an  amount  sufficient  to  generate 
100  1.  of  acetylene  measured  at  20°  and  740  mm.  ? 

18.  How  would  the  volume  of  a  definite  amount  of  carbon  monox- 
ide compare  with  the  volume  of  carbon  dioxide  formed  by  its  com- 
bustion, the  measurements  being  made  under  the  same  conditions? 


CHAPTER   XVIII 
FLAMES,  —  ILLUMINANTS 

Conditions  necessary  for  flames.  It  has  been  seen  that 
when  two  substances  unite  chemically,  with  the  production 
of  light  and  heat,  the  act  of  union  is  called^  combustion. 
When  one  of  the  substances  undergoing  combustion  remains 
solid  at  the  temperature  occasioned  by  the  combustion,  light 
may  be  given  off,  but  there  is  no  flame.  Thus  iron  wire 
burning  in  oxygen  throws  off  a  shower  of  sparks  and  is 
brilliantly  incandescent,  but  no  flame  is  seen.  When,  how- 
ever, both  of  the  substances  are  gases  or  vapors  at  the 
temperature  reached  in  the  combustion,  the  act  of  union 
is  accompanied  by  a  flame. 

Flames  from  burning  liquids  or  solids.  Many  substances  which  are 
liquids  or  solids  at  ordinary  temperatures  burn  with  a  flame  because 
the  heat  of  combustion  vaporizes  them  slowly,  and  the  flame  is  due 
to  the  union  of  this  vapor  with  the  gas  supporting  the  combustion. 

Supporter  of  combustion.  That  gas  which  surrounds  the 
flame  and  constitutes  the  atmosphere  in  which  the  com- 
bustion occurs  'is  said  to  support  the  combustion.  The 
other  gas  which  issues  into  this  atmosphere  is  said  to  be 
the  combustible  gas.  Thus,  in  the  ordinary  combustion  of 
coal  gas  in  the  air  the  coal  gas  is  said  to  be  combustible, 
while  the  air  is  regarded  as  the  supporter  of  combustion. 
These  terms  are  entirely  relative,  however,  for  a  jet  of 
air  issuing  into  an  atmosphere  of  coal  gas  will  bur 
when  ignited,  the  coal  gas  supporting  the  combustio 

213 


or 
rl?«^BB" 


214     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


Gas-* 


Ordinarily,  when  we  say  that  a  gas   is   combustible  we 
mean  that  it  is  combustible  in  an  atmosphere  of  air. 

Either  gas  may  be  the  supporter  of  combustion.  That  the  terms 
combustible  and  supporter  of  combustion  are  merely  relative  may 
be  shown  in  the  following  way :  A  lamp  chimney  A  is  fitted  with 
a  cork  and  glass  tubes,  as  shown  in  Fig.  62.  The  tube  C  should 
have  a  diameter  of  from  12  to  15  mm.  A  thin 
sheet  of  asbestos  in  which  is  cut  a  circular  open- 
ing about  2  cm.  in  diameter  is  placed  over  the 
top  of  the  chimney.  The  opening  in  the  asbestos 
is  closed  with  the  palm  of  the  hand,  and  gas  is 
admitted  to  the  chimney  through  the  tube  B. 
The  air  in  the  chimney  is  soon  expelled  through 
the  tube  C,  and  the  gas  itself  is  then  lighted  at 
the  lower  end  of  this  tube.  The  hand  is  now 
removed  from  the  opening  in  the  asbestos,  when 
the  flame  at  the  end  of  the  tube  at  once  rises  and 
appears  at  the  end  within  the  chimney,  as  shown 
in  the  figure.  The  excess  of  coal  gas  now  escapes 
from  the  opening  in  the  asbestos  and  may  be 
lighted.  The  flame  at  the  top  of  the  asbestos  board  is  due  to  the 
combustion  of  coal  gas  in  air,  while  the  flame  within  the  chimney 
is  due  to  the  combustion  of  air  in  coal  gas,  the  air  being  drawn  up 
through  the  tube  by  the  escaping  gas. 

Appearance  of  flames.  The  flame  caused  by  the  union 
of  hydrogen  and  oxygen  is  almost  colorless  and  invisible. 
Chlorine  and  hydrogen  combine  with  a  pale  violet  flame, 
carbon  monoxide  burns  in  oxygen  with  a  blue  flame, 
while  ammonia  burns  with  a  deep  yellow  flame.  The  color 
and  appearance  of  flames  are  therefore  often  quite  charac- 
teristic of  the  particular  combustion  which  occasions  them. 

Structure  of  flames.  When  the  gas  undergoing  combus- 
tion issues  from  a  round  opening  into  an  atmosphere  of 
the  gas  supporting  combustion,  as  is  the  case  with  the 
burning  Bunsen  burner  (Fig.  63),  the  flame  is  generally 


FIG. 62 


FLAMES,  —  ILLUMINANTS  2 1 5 

conical  in  outline.  It  consists  of  several  distinct  cones, 
one  within  the  other,  the  boundary  between  them  being 
marked  by  differences  of  color  or  luminosity.  In  the  sim- 
plest flame,  of  which  hydrogen  burning  in  oxygen  is  a  good 
example,  these  cones  are  two  in  number,  —  an  inner  one, 
formed  by  unburned  gas,  and  an  outer  one,  usually  more 
or  less  luminous,  consisting  of  the  combining 
gases.  This  outer  one  is  in  turn  surrounded 
by  a  third  envelope  of  the  products  of  combus- 
tion; this  envelope  is  sometimes  invisible,  as 
in  the  present  case,  but  is  sometimes  faintly 
luminous.  The  lower  part  of  the  inner  cone  of 
the  flame  is  quite  cool  and  consists  of  unburned  .,.„.. 
gas.  Toward  the  top  of  the  inner  cone  the  gas 
has  become  heated  to  a  high  temperature  by 
the  burning  envelope  surrounding  it.  On  reach- 
ing the  supporter  of  combustion  on  the  outside 
it  is  far  above  its  kindling  temperature,  and  FlG-63 
combustion  follows  with  the  evolution  of  much  heat.  The 
region  of  combustion  just  outside  the  inner  cone  is  there- 
fore the  hottest  part  of  the  flame. 

Oxidizing  and  reducing  flames.  Since  the  tip  of  the  out- 
side cone  consists  of  very  hot  products  of  combustion 
mixed  with  oxygen  from  the  air,  a  substance  capable  of 
oxidation  placed  in  this  part  of  the  flame  becomes  very 
hot  and  is  easily  oxidized.  The  oxygen  with  which  it  com- 
bines comes,  of  course,  from  the  atmosphere,  and  not  from 
the  products  of  combustion.  This  outer  tip  of  the  flame  is 
called  the  oxidizing  flame . 

At  the  tip  of  the  inner  cone  the  conditions  are  quite 
different.  This  region  consists  of  a  highly  heated  com- 
bustible gas,  which  has  not  yet  reached  a  supply  of  oxygen. 


216     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

If  a  substance  rich  in  oxygen,  such  as  a  metallic  oxide,  is 
placed  in  this  region  of  the  flame,  the  heated  gases  com- 
bine with  its  oxygen  and  the  substance  is  reduced. 
This  part  of  the  flame  is  called  the  reducing  flame. 
These  flames  are  used  in  testing  certain  sub- 
stances, especially  minerals.  For  this  purpose 
they  are  produced  by  blowing  into  a  small  lumi- 
nous Bunsen  flame  from  one  side  through  a  blow- 
pipe. This  is  a  tube  of  the  *^ 
shape  shown  in  Fig.  64. 
FIG.  64  The  flame  is  Directed  in 

any  desired  way  and  has  the  oxi- 
dizing and  reducing  regions  very 
clearly  marked  (Fig.  65).  It  is 

non-luminous  from  the  same  causes  which  render  the  open 
Bunsen  burner  flame  non-luminous,  the  gases  from  the 
lungs  serving  to  furnish  oxygen  and  to  dilute  the  com- 
bustible gas. 

Luminosity  of  flames.  The  luminosity  of  flames  is  due  to  a  num- 
ber of  distinct  causes,  and  may  therefore  be  increased  or  diminished 
in  several  ways. 

i.  Presence  of  solid  matter.  The  most  obvious  of  these  causes 
is  the  presence  in  the  flame  of  incandescent  solid  matter.  Thus 
chalk  dust  sifted  into  a  non-luminous  flame  renders  it  luminous. 
When  hydrocarbons  form  a  part  of  the  combustible  gas,  as  they  do 
in  nearly  all  illuminating  gases  and  oils,  some  carbon  is  usually  set 
free  in  the  process  of  combustion.  This  is  made  very  hot  by  the 
flame  and  becomes  incandescent,  giving  out  light.  In  a  well-regu- 
lated flame  it  is  afterward  burned  up,  but  when  the  supply  of  oxygen 
is  insufficient  it  escapes  from  the  flame  as  lampblack  or  soot.  That 
it  is  temporarily  present  in  a  well-burning  luminous  flame  may  be 
demonstrated  by  holding  a  cold  object,  such  as  a  small  evaporating 
dish,  in  the  flame  for  a  few  seconds.  This  cold  object  cools  the 
carbon  below  its  kindling  temperature,  and  it  is  deposited  on  the 
object  as  soot. 


FLAMES,  —  ILLUMINANTS  2  1 7 

2.  Pressure.    A  second  factor  in  the  luminosity  of  flames  is  the 
pressure    under    which    the    gases    are   burning.     Under    increased 
pressure  there  is  more  matter  in  a  given  volume  of  a  gas,  and  the 
chemical  action  is  more  energetic  than  when  the  gases  are  rarefied. 
Consequently  there  is  more  heat  and  light.    A  candle  burning  on  a 
high  mountain  gives  less  light  than  when  it  burns  at  the  sea  level. 

.If  the  gas  is  diluted  with  a  non-combustible  gas,  the  effect  is 
the  same  as  if  it  is  rarefied,  for  under  these  conditions  there  is  less 
combustible  gas  in  a  given  volume. 

3.  Temperature.    The  luminosity  also  depends  upon  the  tempera- 
ture attained  in  the  combustion.    In  general  the  hotter  the  flame  the 
greater  the  luminosity ;  hence  cooling  the  gases  before  combustion 
diminishes  the  luminosity  of  the  flame  they  will  make,  because  it 
diminishes  the  temperature  attained  in  the  combustion.    Thus  the 
luminosity  of  the  Bunsen   flame  is  largely  diminished  by  the  air 
drawn  up  with  the  gas.    This  is  due  in  part  to  the  fact  that  the 
burning  gas  is  diluted  and  cooled  by  the  air  drawn  in.    The  oxygen 
thus  introduced  into  the  flame  also  causes  the  combustion  of  the  hot 
particles  of  carbon  which  would  otherwise  tend  to  make  the  flame 
luminous. 

Illuminating  and  fuel  gases.  A  number  of  mixtures  of 
combustible  gases,  consisting  largely  of  carbon  compounds 
and  hydrogen,  find  extensive  use  for  the  production  of 
light  and  heat.  The  three  chief  varieties  are  coal  gas, 
water  gas,  and  natural  gas.  The  use  of  acetylene  gas 
has  already  been  referred  to. 

Coal  gas.  Coal  gas  is  made  by  heating  bituminous  coal 
in  large  retorts  out  of  contact  with  the.air.  Soft  or  bitumi- 
nous coal  contains,  in  addition  to  large  amounts  of  carbon, 
considerable  quantities  of  compounds  of  hydrogen,  oxygen, 
nitrogen,  and  sulphur.  When  distilled  the  nitrogen  is  lib- 
erated partly  in  the  form  of  ammonia  and  cyanides  and 
partly  as  free  nitrogen  gas ;  the  sulphur  is  converted  into 
hydrogen  sulphide,  carbon  disulphide,  and  oxides  of  sul- 
phur ;  the  oxygen  into  water  and  oxides  of  carbon.  The 


218     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


remaining  hydrogen  is  set  free  partly  as  hydrogen  and 
partly  in  combination  with  carbon  in  the  form  of  hydro- 
carbons. The  most  important  of  these  is  methane,  with 
smaller  quantities  of  many  others,  some  of  which  are  liquids 
or  solids  at  ordinary  temperatures.  The  great  bulk  of  the 
carbon  remains  behind  as  coke  and  retort  carbon. 

The  manufacture  of  coal  gas.  In  the  manufacture  of  coal  gas  it 
is  necessary  to  separate  from  the  volatile  constituents  formed  by  the 
heating  of  the  coal  all  those  substances  which  are  either  solid  or 
liquid  at  ordinary  temperature,  since  these  would  clog  the  gas  pipes. 

Certain  gaseous 
constituents,  such 
as  hydrogen  sul- 
phide and  ammo- 
nia, must  also  be 
removed.  The 
method  used  to 
accomplish  this  is 
shown  in  Fig.  66. 
The  coal  is  heated 
in  air-tight  retorts 
illustrated  by  A . 

The  volatile  products  escape  through  the  pipe  X  and  bubble  into  the 
tarry  liquid  in  the  large  pipe  B,  known  as  the  hydraulic  main,  which 
runs  at  right  angles  to  the  retorts.  Here  is  deposited  the  greater  por- 
tion of  the  solid  and  liquid  products,  forming  a  tarry  mass  known 
as  coal  tar.  Much  of  the  ammonia  also  remains  dissolved  in  this 
liquid.  The  partially  purified  gas  then  passes  into  the  pipes  C,  which 
serve  to  cool  it  and  further  remove  the  solid  and  liquid  matter.  The 
gas  then  passes  into  D,  which  is  filled  with  coke  over  which  a  jet  of 
water  is  sprayed.  The  water  still  further  cools  the  gas  and  at  the 
same  time  partially  removes  such  gaseous  products  as  hydrogen  sul- 
phide and  ammonia,  which  are  soluble  in  water.  In  E  the  gas  passes 
over  some  material  such  as  lime,  which  removes  the  last  portions  of 
the  sulphur  compounds  as  well  as  much  of  the  carbon  dioxide  present. 
From  E  the  gas  passes  into  the  large  gas  holder  F,  from  which  it  is 
distributed  through  pipes  to  the  places  where  it  is  burned. 


FIG.  66 


FLAMES,  —  ILLUMINANTS  2  1  9 

One  ton  of  good  gas  coal  yields  approximately  10,000  cu.  ft.  of 
gas,  1400  Ib.  of  coke,  120  Ib.  of  tar,  and  20  gal.  of  ammoniacal 
liquor. 

Not  only  is  the  ammonia  obtained  in  the  manufacture  of  the  gas  of 
great  importance,  but  the  coal  tar  also  serves  as  the  source  of  many 
very  useful  substances,  as  will  be  explained  in  Chapter  XXXII. 

Water  gas.  Water  gas  is  essentially  a  mixture  of  carbon 
monoxide  and  hydrogen.  It  is  made  by  passing  steam 
over  very  hot  anthracite  coal,  when  the  reaction  shown  in 
the  following  equation  takes  place  : 


When  required  merely  to  produce  heat"  the  gas  is  at  once 
ready  for  use.  When  made  for  illuminating  purposes  it 
must  be  enriched,  that  is,  illuminants  must  be  added,  since 
both  carbon  monoxide  and  hydrogen  burn  with  non-lumi- 
nous flames.  This  is  accomplished  by  passing  it  into  heaters 
containing  highly  heated  petroleum  oils.  The  gas  takes 
up  hydrocarbon  gases  formed  in  the  decomposition  of  the 
petroleum  oils,  which  make  it  burn  with  a  luminous  flame. 

Water  gas  is  very  effective  as  a  fuel,  since  both  carbon 
monoxide  and  hydrogen  burn  with  very  hot  flames.  It  has 
little  odor  and  is  very  poisonous.  Its  use  is  therefore 
attended  with  some  risk,  since  leaks  in  pipes  are  very  likely 
to  escape  notice. 

Natural  gas.  This  substance,  so  abundant  in  many  locali- 
ties, varies  much  in  composition,  but  is  composed  princi- 
pally of  methane.  When  used  for  lighting  purposes  it  is 
usually  burned  in  a  burner  resembling  an  open  Bunsen,  the 
illumination  being  furnished  by  an  incandescent  mantle. 
This  is  the  case  in  the  familiar  Welsbach  burner.  Con- 
trary to  statements  frequently  made,  natural  gas  contains 
no  free  hydrogen. 


220     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


TABLE  SHOWING  COMPOSITION  OF  GASES 


PENNSYLVANIA 
NATURAL 
GAS 

COAL 
GAS 

WATER 
GAS 

ENRICHED 
WATER 
GAS 

Hydrogen        .... 

A  I   7 

52  88 

•2Q    OO 

Methane 

QO  0)d. 

At.  6 

2  16 

24.  oo 

12  OC 

Carbon  monoxide     .     .     . 

6.4 

36  80 

2O  OO 

Carbon  dioxide    .     . 
Nitrogen 

0.30 

906 

2.0 
I  2 

3-47 

4f)(\ 

0.30 
•7    Cn 

Oxygen  . 

.WVJ 

Q.T. 

•uy 

~.}<~> 

i  so 

Hydrocarbon  vapors 

.    .    .    . 

i-5 

•     •     v 

1.50 

These  are  analyses  of  actual  samples,  and  may  be  taken  as  about 
the  average  for  the  various  kinds  of  gases.  Any  one  of  these  may 
vary  considerably.  The  nitrogen  and  oxygen  in  most  cases  is  due 
to  a  slight  admixture  of  air  which  is  difficult  to  exclude  entirely  in 
the  manufacture  and  handling  of  gases. 

Fuels.  A  variety  of  substances  are  used  as  fuels,  the 
most  important  of  them  being  wood,  coal,  and  the  various 
gases  mentioned  above.  Wood  consists  mainly  of  com- 
pounds of  carbon,  hydrogen,  and  oxygen.  The  composi- 
tion of  coal  and  the  fuel  gases  has  been  given.  Since 
these  fuels  are  composed  principally  of  carbon  and  hydro- 
gen or  their  compounds,  the  chief  products  of  combustion 
are  carbon  dioxide  and  water.  The  practice  of  heating 
rooms  with  portable  gas  or  oil  stoves  with  no  provision 
for  removing  the  products  of  combustion  is  to  be  con- 
demned, since  the  carbon  dioxide  is  generated  in  sufficient 
quantities  to  render  the  air  unfit  for  breathing.  Rooms 
so  heated  also  become  very  damp  from  the  large  amount 
of  water  vapor  formed  in  the  combustion,  and  which  in 


FLAMES,  —  ILLUMINANTS 


221 


cold  weather  condenses  on  the  window  glass,  causing  the 
glass  to  "  sweat."  Both  coal  and  wood  contain  a  certain 
amount  of  mineral  substances  which  constitute  the  ashes. 

The  electric  furnace.  In  recent  years  electric  furnaces  have  come 
into  wide  use  in  operations  requiring  a  very  high  temperature.  Tem- 
peratures as  high  as  3500°  can  be  easily 
reached,  whereas  the  hottest  oxyhydro- 
gen  flame  is  not  much  above  2000°. 
These  furnaces  are  constructed  on  one 
of  two  general  principles. 

i.  Arc  furnaces.  In  the  one  type 
the  source  of  heat  is  an  electric  arc 
formed  between  carbon  electrodes  sepa- 
rated a  little  from  each  other,  as  shown 


FIG. 67 


in  Fig.  67.  The  substance  to  be  heated  is  placed  in  a  vessel,  usually 
a  graphite  crucible,  just  below  the  arc.  The  electrodes  and  crucible 
are  surrounded  by  materials  which  fuse  with  great  difficulty,  such  as 
magnesium  oxide,  the  walls  of  the  furnace  being  so  shaped  as  to  reflect 
the  heat  downwards  upon  the  contents  of  the  crucible. 

2.  Resistance  furnaces.  In  the  other  type  of  furnace  the  heat  is 
generated  by  the  resistance  offered  to  the  current  in  its  passage 
through  the  furnace.  In  its  simplest  form  it  may  be  represented  by 

Fig.  68.  The  furnace 
is  merely  a  rectangular 
box  built  up  of  loose 
bricks.  The  elec- 
trodes E,  each  consist- 
ing of  a  bundle  of 
FlG  68  carbon  rods,  are  intro- 

duced   through    the 

sides  of  the  furnace.  The  materials  to  be  heated,  C,  are  filled  into 
the  furnace  up  to  the  electrodes,  and  a  layer  of  broken  coke  is 
arranged  so  as  to  extend  from  one  electrode  to  the  other.  More  of 
the  charge  is  then  placed  on  top  of  the  coke.  In  passing  through  the 
broken  coke  the  electrical  current  encounters  great  resistance.  This 
generates  great  heat,  and  the  charge  surrounding  the  coke  is  brought 
to  a  very  high  temperature.  The  advantage  of  this  type  of  furnace  is 
that  the  temperature  can  be  regulated  to  any  desired  intensity. 


222     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


EXERCISES 

1.  Why  does  charcoal  usually  burn  with  no  flame  ?    How  do  you 
account  for  the  flame  sometimes  observed  when  it  burns? 

2.  How  do  you*  account  for  the  fact  that  a  candle  burns  with  a 
flame  ? 

3.  What  two  properties  must  the  mantle  used  in  the  Welsbach 
lamp  possess  ? 

4.  (#)  In  what  respects  does  the  use  of  the  Welsbach  mantle 
resemble  that  of  lime  in  the  calcium  light?    (£)  If  the  mantle  were 
made  of  carbon,  would  it  serve  the  same  purpose  ? 

5.  Would  anthracite  coal  be  suitable  for  the  manufacture  of 
coal  gas? 

6.  How  could  you  prove  the  formation  of  carbon  dioxide  and 
water  in  the  combustion  of  illuminating  gases  ? 

7.  Suggest  a  probable  way  in  which  natural  gas  has  been  formed. 

8.  Coal  frequently  contains  a  sulphide  of  iron,     (a)  What  two 
sulphur  compounds  are  likely  to  be  formed  when  gas  is  made  from 
such  coal?  (^)  Suggest  some  suitable  method  for  the  removal  of 
these  compounds. 

9.  Why  does  the  use  of  the  bellows  on  the  blacksmith's  forge 
cause  a  more  intense  heat  ? 

10.  What  volume  of  oxygen  is  necessary  to  burn  100  1.  of  marsh 
gas  and  what  volume  of  carbon  dioxide  would  be  formed,  all  of  the 
gases  being  measured  under  standard  conditions  ? 

11.  Suppose  a  cubic  meter  of  Pennsylvania  natural  gas,  measured 
under  standard  conditions,  were  to  be  burned.     How  much  water  by 
weight  would  result  ? 


CHAPTER    XIX 
MOLECULAR  WEIGHTS,   ATOMIC   WEIGHTS,   FORMULAS 

Introduction.  In  the  chapter  on  The  Atomic  Theory,  it 
was  shown  that  if  it  were  true  that  two  elements  uniting  to 
form  a  compound  always  combined  in  the  ratio  of  one  atom 
of  one  element  to  one  atom  of  the  other  element,  it  would 
be  a  very  easy  matter  to  decide  upon  figures  which  would 
represent  the  relative  weights  of  the  different  atoms.  It 
would  only  be  necessary  to  select  some  one  element  as  a 
standard  and  determine  the  weight  of  every  element  which 
combines  with  a  definite  weight  (say  I  g.)  of  the  standard 
element.  The  figures  so  obtained  would  evidently  repre- 
sent the  relative  weights  of  the  atoms. 

But  the  law  of  multiple  proportion  at  once  reminds  us 
that  two  elements  may  unite  in  several  proportions ;  and 
there  is  no  simple  way  to  determine  the  number  of  atoms 
present  in  the  molecule  of  any  compound.  Consequently 
the  problem  of  deciding  upon  the  relative  atomic  weights 
is  not  an  easy  one.  To 'the  solution  of  this  problem  we 
must  now  turn. 

Dalton's  method  of  determining  atomic  weights.  When  Dalton  first 
advanced  the  atomic  theory  he  attempted  to  solve  this  problem  by 
very  simple  methods.  He  thought  that  when  only  one  compound 
of  two  elements  is  known  it  is  reasonable  to  suppose  that  it  contains 
one  atom  of  each  element.  He  therefore  gave  the  formula  HO  to 
water,  and  HN  to  ammonia.  When  more  than  two  compounds  were 
known  he  assumed  that  the  most  familiar  or  the  most  stable  one 
had  the  simple  formula.  He  then  determined  the  atomic  weight  as 

223 


224     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

explained  above.  The  results  he  obtained  were  contradictory  and 
very  far  from  satisfactory,  and  it  was  soon  seen  that  some  other 
method,  resting  on  much  more  scientific  grounds,  must  be  found  to 
decide  what  compounds,  if  any,  have  a  single  atom  of  each  element 
present. 

Determination  of  atomic  weights.  Three  distinct  steps 
are  involved  in  the  determination  of  the  atomic  weight  of 
an  element:  (i)  determination  of  the  equivalent,  (2)  deter- 
mination of  molecular  weights  of  its  compounds,  and  (3) 
deduction  of  the  exact  atomic  weight  from  the  equivalent 
and  molecular  weights. 

i.  Determination  of  the  equivalent.  By  the  equivalent 
of  an  element  is  meant  the  weight  of  the  element  which 
will  combine  with  a  fixed  weight  of  some  other  element 
chosen  as  a  standard.  It  has  already  been  explained  that 
oxygen  has  been  selected  as  the  standard  element  for 
atomic  weights,  with  a  weight  of  16.  This  same  standard 
will  serve  very  well  as  a  standard  for  equivalents.  The 
equivalent  of  an  element  is  the  weight  of  the  element  which 
will  combine  ^vith  id  g.  of  oxygen.  Thus  16  g.  of  oxygen 
combines  with  16.03  g-  °f  sulphur,  65 .4  g.  of  zinc,  2 1 5.86  g. 
of  silver,  70.9  g.  of  chlorine.  These  figures,  therefore, 
represent  the  equivalent  weights  of  these  elements. 

Relation  of  atomic  weights  to  equivalents.  According 
to  the  atomic  theory  combination  always  takes  place 
between  whole  numbers  of  atoms.  Thus  one  atom  unites 
with  one  other,  or  with  two  or  three ;  or  two  atoms  may 
unite  with  three,  or  three  with  five,  and  so  on. 

When  oxygen  combines  with  zinc  the  combination  must 
be  between  definite  numbers  of  the  two  kinds  of  atoms. 
Experiment  shows  that  these  two  elements  combine  in  the 
ratio  of  16  g.  of  oxygen  to  65.4  g.  of  zinc.  If  one  atom  of 


MOLECULAR  AND  ATOMIC  WEIGHTS          225 

oxygen  combines  with  one  atom  of  zinc,  then  this  ratio  must 
be  the  ratio  between  the  weights  of  the  two  atoms.  If  one 
atom  of  oxygen  combines  with  two  atoms  of  zinc,  then  the 
ratio  between  the  weights  of  the  two  atoms  will  be  16  :  32.7. 
If  two  atoms  of  oxygen  combine  with  one  atom  of  zinc, 
the  ratio  by  weight  between  the  two  atoms  will  be  8  :  65.4. 
It  is  evident,  therefore,  that  the  real  atomic  weight  of  an 
element  must  be  some  multiple  or  submultiple  of  the  equiv- 
alent;  in  other  words,  the  equivalent  multiplied  by  J-,  I, 
2,  or  3  will  give  the  atomic  weight. 

Combining  weights.  A  very  interesting  relation  holds  good  between 
the  equivalents  of  the  various  elements.  We  have  just  seen  that  the 
figures  16.03,  654,  215.86,  and  70.9  are  the  equivalents  respectively 
of  sulphur,  zinc,  silver,  and  chlorine.  These  same  figures  represent 
the  ratios  by  weight  in  which  these  elements  combine  among  them- 
selves. Thus  215.86  g.  of  silver  combine  with  70.9  g.  of  chlorine 
and  with  2  x  16.03  £•  °f  sulphur.  65.4  g.  of  zinc  combine  with 
70.9  g.  of  chlorine  and  2  x  16.03  g-  of  .sulphur. 

By  taking  the  equivalent  or  some  multiple  of  it  a  value  can  be 
obtained  for  each  element  which  will  represent  its  combining  value, 
and  for  this  reason  is  called  its  combining  weight.  It  is  important 
to  notice  that  the  fact  that  a  combining  weight  can  be  obtained  for 
each  element  is  not  a  part  of  a  theory,  but  is  the  direct  result  of 
experiment. 

Elements  with  more  than  one  equivalent.  It  will  be 
remembered  that  oxygen  combines  with  hydrogen  in  two 
ratios.  In  one  case  16  g.  of  oxygen  combine  with  2.016  g. 
of  hydrogen  to  form  water;  in  the  other  16  g.  of  oxygen 
combine  with  1.008  g.  of  hydrogen  to  form  hydrogen  diox- 
ide. The  equivalents  of  hydrogen  are  therefore  2.016  and 
1.008.  Barium  combines  with  oxygen  in  two  proportions  : 
in  barium  oxide  the  proportion  is  16  g.  of  oxygen  to  137.4  g. 
of  barium ;  in  barium  dioxide  the  proportion  is  16  g.  of 
oxygen  to  68.7  g.  of  barium. 


226     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

In  each  case  one  equivalent  is  a  simple  multiple  of  the 
other,  so  the  fact  that  there  may  be  two  equivalents  does 
not  add  to  the  uncertainty.  All  we  knew  before  was  that 
the  true  atomic  weight  is  some  multiple  of  the  equivalent. 

2.  The  determination  of  molecular  weights.  To  decide 
the  question  as  to  which  multiple  of  the  equivalent  cor- 
rectly represents  the  atomic  weight  of  an  element,  it  has 
been  found  necessary  to  devise  a  method  of  determining 
the  molecular  weights  of  compounds  containing  the  element 
in  question.  Since  the  molecular  weight  of  a  compound  is 
merely  the  sum  of  the  weights  of  all  the  atoms  present  in 
it,  it  would  seem  to  be  impossible  to  determine  the  molec- 
ular weight  of  a  compound  without  first  knowing  the 
atomic  weights  of  the  constituent  atoms,  and  how  many 
atoms  of  each  element  are  present  in  the  molecule.  But 
certain  facts  have  been  discovered  which  suggest  a  way  in 
which  this  can  be  done. 

Avogadro's  hypothesis.  We  have  seen  that  the  laws  of 
Boyle,  Charles,  and  Gay-Lussac  apply  to  all  gases  irre- 
spective of  their  chemical  character.  This  would  lead  to 
the  inference  that  the  structure  of  gases  must  be  quite 
simple,  and  that  it  is  much  the  same  in  all  gases. 

In  1811  Avogadro,  an  Italian  physicist,  suggested  that 
if  we  assume  all  gases  under  the  same  conditions  of 
temperature  and  pressure  to  have  the  same  number  of 
molecules  in  a  given  volume,  we  shall  have  a  probable 
explanation  of  the  simplicity  of  the  gas  laws.  It  is  diffi- 
cult to  prove  the  truth  of  this  hypothesis  by  a  simple 
experiment,  but  there  are  so  many  facts  known  which 
are  in  complete  harmony  with  this  suggestion  that  there 
is  little  doubt  that  it  expresses  the  truth.  Avogadro's 
hypothesis  may  be  stated  thus :  Equal  volumes  of  all 


MOLECULAR  AND  ATOMIC  WEIGHTS          227 

gases  under  the  same  conditions  of  temperature  and  pres- 
sure contain  the  same  number  of  molecules. 

Avogadro's  hypothesis  and  molecular  weights.  Assum- 
ing that  Avogadro's  hypothesis  is  correct,  we  have  a  very 
simple  means  for  deciding  upon  the  relative  weights  of 
molecules ;  for  if  equal  volumes  of  two  gases  contain  the 
same  number  of  molecules,  the  weights  of  the  two  volumes 
must  be  in  the  same  ratio  as  the  weights  of  the  individual 
molecules  which  they  contain.  If  we  adopt  some  one  gas 
as  a  standard,  we  can  express  the  weights  of  all  other  gases 
as  compared  with  this  one,  and  the  same  figures  will  express 
the  relative  weights  of  the  molecules  of  which  the  gases 
are  composed. 

Oxygen  as  the  standard.  It  is  important  that  the  same 
standard  should  be  adopted  for  the  determination  of  molec- 
ular weights  as  has  been  decided  upon  for  atomic  weights 
and  equivalents,  so  that  the  three  values  may  be  in  harmony 
with  each  other.  Accordingly  it  is  best  to  adopt  oxygen 
as  the  standard  element  with  which  to  compare  the  molec- 
ular weights  of  other  gases,  being  careful  to  keep  the 
oxygen  atom  equal  to  16. 

The  oxygen  molecule  contains  two  atoms.  One  point 
must  not  be  overlooked,  however.  We  desire  to  have  our 
unit,  the  oxygen  atom,  equal  to  16.  The  method  of  com- 
paring the  weights  of  gases  just  suggested  compares  the 
molecules  of  the  gases  with  the  molecule  of  oxygen.  Is  the 
molecule  and  the  atom  of  oxygen  the  same  thing  ?  This 
question  is  answered  by  the  following  considerations. 

We  have  seen  that  when  steam  is  formed  by  the  union 
of  oxygen  and  hydrogen,  two  volumes  of  hydrogen  combine 
with  one  volume  of  oxygen  to  form  two  volumes  of  steam. 
Let  us  suppose  that  the  one  volume  of  oxygen  contains  100 


228     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

molecules ;  then  the  two  volumes  of  steam  must,  according 
to  Avogadro's  hypothesis,  contain  200  molecules.  But 
each  of  these  200  molecules  must  contain  at  least  one  atom 
of  oxygen,  or  200  in  all,  and  these  200  atoms  came  from 
100  molecules  of  oxygen.  It  follows  that  each  molecule  of 
oxygen  must  contain  at  least  two  atoms  of  oxygen. 

Evidently  this  reasoning  merely  shows  that  there  are 
at  least  two  atoms  in  the  oxygen  molecule.  There  may  be 
more  than  that,  but  as  there  is  no  evidence  to  this  effect, 
we  assume  that  the  molecule  contains  two  atoms  only. 

It  is  evident  that  if  we  wish  to  retain  the  value  16  for 
the  atom  of  oxygen  we  must  take  twice  this  value,  or  32, 
for  the  value  of  the  oxygen  molecule,  when  using  it  as  a 
standard  for  molecular  weights. 

Determination  of  the  molecular  weights  of  gases  from 
their  weights  compared  with  oxygen.  Assuming  the  molec- 
ular weight  of  oxygen  to  be  32,  Avogadro's  hypothesis 
gives  us  a  ready  means  for  determining  the  molecular 
weight  of  any  other  gas,  for  all  that  is  required  is  to  know 
its  weight  compared  with  that  of  an  equal  volume  of  oxygen. 
For  example,  I  1.  of  chlorine  is  found  by  experiment  to 
weigh  2.216  times  as  much  as  1 1.  of  oxygen.  The  molecular 
weight  of  chlorine  must  therefore  be  2.216x32,  or  70.91. 

If,  instead  of  comparing  the  relative  weights  of  i  1.  of 
the  two  gases,  we  select  such  a  volume  of  oxygen  as  will 
weigh  32  g.,  or  the  weight  in  grams  corresponding  to  the 
molecular  weight  of  the  gas,  the  calculation  is  much  sim- 
plified. It  has  been  found  that  32  g.  of  oxygen,  under 
standard  conditions,  measure  22.4  1.  This  same  volume 
of  hydrogen  weighs  2^019  g. ;  of  chlorine  70.9  g. ;  of  hydro- 
chloric acid  36.458  g.  The  weights  of  these  equal  volumes 
ntust  be  proportional  to  their  molecular  weights,  and  since 


MOLECULAR  AND  ATOMIC  WEIGHTS 


229 


the  weight  of  the  oxygen  is  the  same  as  the  value  of  its 
molecular  weight,  so  too  will  the  weights  of  the  22.4  1.  of 
the  other  gases  be  equal  to  the  value  of  their  molecular 
weights. 

As  a  summary  we  can  then  make  the  following  state- 
ment :  The  molecular  weight  of  any  gas  may  be  determined 
by  calculating  the  weight  of  22,4  I.  of  the  gas,  meastired 
under  standard  conditions. 

Determination  of  molecular  weights  from  density  of  gases. 
In  an  actual  experiment  it  is  easier  to  determine  the 
density  of  a  gas  than  the  weight  of  a  definite  volume  of 
it.  The  density  of  a  gas  is  usually  defined  as  its  weight 
compared  with  that  of  an  equal  volume  of  air.  Having 
determined  the  density  of  a  gas,  its  weight  compared  with 
oxygen  may  be  determined  by  multiplying  its  density  by 
the  ratio  between  the  weights  of  air  and  oxygen.  This 
ratio  is  0.9046.  To  compare  it  with  our  standard  for 
atomic  weights  we  must  further  multiply  it  by  32,  since 
the  standard  is  -j^  the  weight  of  oxygen  molecules.  The 
steps  then  are  these  : 

1.  Determine  the  density  of  the  gas  (its  weight  com- 
pared with  air). 

2.  Multiply  by  0.9046   to   make   the   comparison  with 
oxygen  molecules. 

3.  Multiply  by  32  to  make  the  comparison  with  the  unit 
for  atomic  weights. 

We  have,  then,  the  formula : 

molecular  weight  =  density  X  0.9046  X  32  ; 

or,  still  more  briefly, 

M.  =  D.  x  28.9. 

The  value  found  by  this  method  for  the  determination  of 
molecular  weights  will  of  course  agree  with  those  found 


230 


AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


FIG.  69 


by  calculating  the  weight  of  22.4  1.  of  the  gas,  since  both 
methods  depend  on  the  same  principles. 

Determination  of  densities  of  gases.  The  relative  weights  of  equal 
volumes  of  two  gases  can  be  easily  determined.  The  following  is 
one  of  the  methods  used.  A  small  flask,  such  as 
is  shown  in  Fig.  69,  is  filled  with  one  of  the  gases, 
and  after  the  temperature  and  pressure  have  been 
noted  the  flask  is  sealed  up  and  weighed.  The  tip 
of  the  sealed  end  is  then  broken  off,  the  flask  filled 
with  the  second  gas,  and  its  weight  determined. 
If  the  weight  of  the  empty  flask  is  subtracted 
from  these  two  weighings,  the  relative  weights 
of  the  gases  is  readily  found. 

3.  Deduction  of  atomic  weights  from  molecular  weights 
and  equivalents.  We  have  now  seen  how  the  equivalent  of 
an  element  and  the  molecular  weight  of  compounds  con- 
taining the  element  can  be  obtained.  Let  us  see  how  it  is 
possible  to  decide  which  multiple  of  the  equivalent  really  is 
the  true  atomic  weight.  As  an  example,  let  us  suppose  that 
the  equivalent  of  nitrogen  has  been  found  to  be  7.02  and 
that  it  is  desired  to  obtain  its  atomic  weight.  The  next  step 
is  to  obtain  the  molecular  weights  of  a  large  number  of 
compounds  containing  nitrogen.  The  following  will  serve : 


APPROXIMATE 

PERCENTAGE  OF 

PART  OF 

DENSITY  BY 
EXPERIMENT 

MOLECULAR 
WEIGHT 
(D.  x  28.9) 

NITROGEN  BY 
EXPERIMENT 

MOLECULAR 
WEIGHT  DUE 
TO  NITROGEN 

Nitrogen  gas    .     .  -  ,. 

0.9671 

27-95 

IOO.OO 

27-95 

Nitrous  oxide  .     .     . 

I-527 

44-13 

63.70 

27.11 

Nitric  oxide      .     .     . 

1.0384 

30.00 

46.74 

14.02 

Nitrogen  peroxide 

1.580 

45-66 

30-49 

13.90 

Ammonia     .... 

0.591 

I7-°5 

82.28 

14.03 

Nitric  acid    .... 

2.l8o 

63.06 

22.27 

14.03 

Hydrocyanic  acid 

0.930 

•    26.87 

51.90 

!3-94 

MOLECULAR  AND  ATOMIC  WEIGHTS          231 

Method  of  calculation.  The  densities  of  the  various  gases 
in  the  first  column  of  this  table  are  determined  by  experi- 
ment, and  are  fairly  accurate  but  not  entirely  so.  By  mul- 
tiplying these  densities  by  28.9  the  molecular  weights  of 
the  compounds  as  given  in  the  second  column  are  obtained. 
By  chemical  analysis  it  is  possible  to  determine  the  percent- 
age composition  of  these  substances,  and  the  percentages 
of  nitrogen  in  them  as  determined  by  analysis  are  given  in 
the  third  column.  If  each  of  these  molecular  weights  is 
multiplied  in  turn  by  the  percentage  of  nitrogen  in  the 
compound,  the  product  will  be  the  weight  of  the  nitro- 
gen in  the  molecular  weight  of  the  compound.  This  will 
be  the  sum  of  the  weights  of  the  nitrogen  atoms  in  the 
molecule.  These  values  are  given  in  the  fourth  column  in 
the  table. 

If  a  large  number  of  compounds  containing  nitrogen  are 
studied  in  this  way,  it  is  probable  that  there  will  be  included 
in  the  list  at  least  one  substance  whose  molecule  contains  a 
single  nitrogen  atom.  In  this  case  the  number  in  the  fourth 
column  will  be  the  approximate  atomic  weight  of  nitrogen. 
On  comparing  the  values  for  nitrogen  in  the  table  it  will  be 
seen  that  a  number  which  is  approximately  14  is  the  small- 
est, and  that  the  others  are  multiples  of  this.  These 
compounds  of  higher  value,  therefore,  contain  more  than 
one  nitrogen  atom  in  the  molecule. 

Accurate  determination  of  atomic  weights.  Molecular 
weights  cannot  be  determined  very  accurately,  and  con- 
sequently the  part  in  them  due  to  nitrogen  is  a  little  uncer- 
tain, as  will  be  seen  in  the  table.  All  we  can  tell  by  this 
method  is  that  the  true  weight  is  very  near  14.  The 
equivalent  can  however  be  determined  very  accurately,  and 
we  have  seen  that  it  is  some  multiple  or  submultiple 


232     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

of  the  true  atomic  weight.  Since  molecular-weight  deter- 
minations have  shown  that  in  the  case  of  nitrogen  the 
atomic  weight  is  near  14,  and  we  have  found  the  equiv- 
alent to  be  7.02,  it  is  evident  that  the  true  atomic  weight 
is  twice  the  equivalent,  or  7.02  x  2  =  14.04. 

Summary.  These,  then,  are  the  steps  necessary  to 
establish  the  atomic  weight  of  an  element. 

1.  Determine  the  equivalent  accurately  by  analysis. 

2.  Determine  the  molecular  weight  of  a  large  number 
of  compounds  of  the  element,  and  by  analysis  the  part  of 
the  molecular  weight  due   to  the   element.    The  smallest 
number  so  obtained  will  be  approximately  the  atomic  weight. 

3.  Multiply  the  equivalent  by  the  small  whole  number 
(usually  i,  2,  or  3),  which  will  make  a  number  very  close 
to  the  approximate  atomic  weight.    The  figure  so  obtained 
\y,ill  be  the  true  atomic  weight. 

Molecular  weights  of  the  elements.  It  will  be  noticed 
that  the  molecular  weight  of  nitrogen  obtained  by  multi- 
plying its  density  by  28.9  is  28.08.  'Yet  the  atomic  weight 
of  nitrogen  as  deduced  from  a  study  of  its  gaseous  com- 
pounds is  14.04.  The  simplest  explanation  that  can  be 
given  for  this  is  that  the  gaseous  nitrogen  is  made  up  of 
molecules,  each  of  which  contains  two  atoms.  In  this  respect 
it  resembles  oxygen ;  for  we  have  seen  that  an  entirely 
different  line  of  reasoning  leads  us  to  believe  that  the 
molecule  of  oxygen  contains  two  atoms.  When  we  wish  to 
indicate  molecules  of  these  gases  the  symbols  N2  and  O2 
should  be  used.  When  we  desire  to  merely  show  the 
weights  taking  part  in  a  reaction  this  is  not  necessary. 

The  vapor  densities  of  many  of  the  elements  show 
that,  like  oxygen  and  nitrogen,  their  molecules  consist  of 
two  atoms.  In  other  cases,  particularly  among  the  metals, 


MOLECULAR  AND  ATOMIC  WEIGHTS          233 

the  molecule  and  the  atom  are  identical.    Still  other  ele- 
ments have  four  atoms  in  their  molecules. 

While  oxygen  contains  two  atoms  in  its  molecules,  a 
study  of  ozone  has  led  to  the  conclusion  that  it  has  three. 
The  formation  of  ozone  from  oxygen  can  therefore  be 
represented  by  the  equation 

3  02  =  2  08. 

Other  methods  of  determining  molecular  weights.  It  will 
be  noticed  that  Avogadro's  law  gives  us  a  method  by  which 
we  can  determine  the  relative  weights  of  the  molecules  of 
two  gases  because  it  enables  us  to  tell  when  we  are  deal- 
ing with  an  equal  number  of  the  two  kinds  of  molecules. 
If  by  any  other  means  we  can  get  this  information,  we  can 
make  use  of  the  knowledge  so  gained  to  determine  the 
molecular  weights  of  the  two  substances.  •* 

Raoult's  laws.  Two  laws  have  been  discovered  which 
give  us  just  such  information.  They  are  known  as  Raoult's 
laws,  and  can  be  stated  as  follows  : 

1 .  When  weights  of  substances  which  are  proportional  to 
their  molecular  weights  are  dissolved  in  the  same  weight  of 
solvent,  the  rise  of  the  boiling  point  is  the  same  in  each  case. 

2.  When  weights  of  substances  which  are  proportional  to 
their  molecular  weights  are  dissolved  in  the  same  weight 
of  solvent,  the  lowering  of  the  freezing  point  is  the  same  in 
each  case. 

By  taking  advantage  of  these  laws  it  is  possible  to  deter- 
mine when  two  solutions  contain  the  same  number  of 
molecules  of  two  dissolved  substances,  and  consequently 
the  relative  molecular  weights  of  the  two  substances. 

Law  of  Dulong  and  Petit.  In  1819  Dulong  and  Petit 
discovered  a  very  interesting  relation  between  the  atomic 


234     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

weight  of  an  element  and  its  specific  heat,  which  holds 
true  for  elements  in  the  solid  state.  If  equal  weights  of 
two  solids,  say,  lead  and  silver,  are  heated  through  the  same 
range  of  temperature,  as  from  10°  to  20°,  it  is  found  that 
very  different  amounts  of  heat  are  required.  The  amount 
of  heat  required  to  change  the  temperature  of  a  solid  or  a 
liquid  by  a  definite  amount  compared  with  the  amount 
required  to  change  the  temperature  of  an  equal  weight 
of  water  by  the  same  amount  is  called  its  specific  heat. 
Dulong  and  Petit  discovered  the  following  law  :  The  specific 
heat  of  an  element  in  the  solid  form  multiplied  by  its  atomic 
weight  is  approximately  equal  to  the  constant  6.25.  That  is, 

at.  wt.  X  sp.  ht.  =  6.25. 

Consequently,  6     - 

at.  wt.  = 


sp.  ht. 

This  law  is  not  very  accurate,  but  it  is  often  possible  by 
means  of  it  to  decide  upon  what  multiple  of  the  equiva- 
lent is  the  real  atomic  weight.  Thus  the  specific  heat  of 
iron  is  found  by  experiment  to  be  o.  1 12,  and  its  equivalent 
is  27.95.  6.25-^0.112  =  55.8.  We  see,  therefore,  that 
the  atomic  weight  is  twice  the  equivalent,  or  55.9. 

How  formulas  are  determined.  It  will  be  well  in  con- 
nection with  molecular  weights  to  consider  how  the  for- 
mula of  a  compound  is  decided  upon,  for  the  two  subjects 
are  very  closely  associated.  Some  examples  will  make  clear 
the  method  followed. 

The  molecular  weight  of  a  substance  containing  hydro- 
gen and  chlorine  was  36.4.  By  analysis  36.4  parts  of  the 
substance  was  found  to  contain  I  part  of  hydrogen  and 
35.4  parts  of  chlorine.  As  these  are  the  simple  atomic 


MOLECULAR  AND  ATOMIC  WEIGHTS          235 

weights  of  the  two  elements,  the  formula  of  the  compound 
must  be  HC1. 

A  substance  consisting  of  oxygen  and  hydrogen  was 
found  to  have  a  molecular  weight  of  34.  Analysis  showed 
that  in  34  parts  of  the  substance  there  were  2  parts  of 
hydrogen  and  32  parts  of  oxygen.  Dividing  these  figures 
by  the  atomic  weights  of  the  two  elements,  we  get 
2  -f-  i  =  2  for  H;  32-^16  =  2  for  O.  The  formula  is 
therefore  H2O2. 

A  substance  containing  2.04%  H,  32.6%  S,  and  65.3% 
O  was  found  to  have  a  molecular  weight  of  98.  In  these 
98  parts  of  the  substance  there  are  98  X  2.04%  =  2  parts 
of  H,  98  X  32.6%  =  32  parts  of  S,  and  98  X  65.3%  =  64 
parts  of  O.  If  the  molecule  weighs  98,  the  hydrogen  atoms 
present  must  together  weigh  2,  the  sulphur  atoms  32,  and 
the  oxygen  atoms  64.  Dividing  these  figures  by  the  respec- 
tive atomic  weights  of  the  three  elements,  we  have,  for  H, 
2-7-1  =  2  atoms  ;  for  S,  32-^32=1  atom ;  for  O,  64  -4-  16 
'=•  4  atoms.  Hence  the  formula  is  H2SO4. 

We  have,  then,  this  general  procedure  :  Find  the  per- 
centage composition  of  the  substance  and  also  its  molecu- 
lar weight.  Multiply  the  molecular  weight  successively 
by  the  percentage  of  each  element  present,  to  find  the 
amount  of  the  element  in  the  molecular  weight  of  the 
compound.  The  figures  so  obtained  will  be  the  respective 
parts  of  the  molecular  weight  due  to  the  several  atoms. 
Divide  by  the  atomic  weights  of  the  respective  elements, 
and  the  quotient  will  be  the  number  of  atoms  present. 

Avogadro's  hypothesis  and  chemical  calculations.  This 
law  simplifies  many  chemical  calculations. 

I.  Application  to  volume  relations  in  gaseous  reactions. 
Since  equal  volumes  of  gases  contain  an  equal  number  of 


236     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

molecules,  it  follows  that  when  an  equal  number  of  gaseous 
molecules  of  two  or  more  gases  take  part  in  a  reaction,  the 
reaction  will  involve  equal  volumes  of  the  gases.  In  the 
equation  C2H2O4  =  H2O  +  CO2  +  CO, 

since  i  molecule  of  each  of  the  gases  CO2  and  CO  is  set 
free  from  each  molecule  of  oxalic  acid,  the  two  substances 
must  always  be  set  free  in  equal  volumes. 

Acetylene  burns  in  accordance  with  the  equation 

2  C2H2  +  5  O2  -  4  CO2  +  2  H2O. 

Hence  2  volumes  of  acetylene  will  react  with  5  volumes  of 
oxygen  to  form  4  volumes  of  carbon  dioxide  and  2  volumes 
of  steam.  That  the  volume  relations  may  be  correct  a 
gaseous  element  must  be  given  its  molecular  formula. 
Thus  oxygen  must  be  written  O2  and  not  2  O. 

2.  Application  to  weights  of  gases.  It  will  be  recalled 
that  the  molecular  weight  of  a  gas  is  determined  by  ascer- 
taining the  weight  of  22.4  1.  of  the  gas.  This  weight 
in  grams  is  called  the  gram-molecular  weight  of  a  gas. 
If  the  molecular  weight  of  any  gas  is  known,  the  weight  of 
a  liter  of  the  gas  under  standard  conditions  may  be  deter- 
mined by  dividing  its  gram-molecular  weight  by  22.4. 
Thus  the  gram-molecular  weight  of  a  hydrochloric  acid 
gas  is  36.458.  A  liter  of  the  gas  will  therefore  weigh 
36.458-^-22.4  =  1.627  g. 

EXERCISES 

1.  From  the  following  data  calculate  the  atomic  weight  of  sulphur. 
The  equivalent,  as  obtained  by  an  analysis  of  sulphur  dioxide,  is 
15.915.  The  densities  and  compositions  of  a  number  of  compounds 
containing  sulphur  are  as  follows  : 


MOLECULAR  AND  ATOMIC  WEIGHTS 


237 


NAME 


DENSITY 


COMPOSITION  BY  PERCENTAGE 


Hydrosulphuric  acid  .     1.1791     8  =  94.11      H  =    5.89 

Sulphur  dioxide 

Sulphur  trioxide 

Sulphur  chloride 

Sulphuryl  chloride . 

Carbon  disulphide  . 


2.222 

S 

=  50.05 

o 

= 

49-95 

2.74 

S 

=  40.05 

0 

= 

59-95 

470 

S 

=  47.48 

Cl 

= 

52.52 

4.64 

S 

=  23.89 

Cl 

= 

5243 

2.68 

S 

=  84.24 

c 

= 

15.76 

O  =  23.89 


2.    Calculate    the    formulas    for    compounds    of    the    following 


compositions : 

(1)  S    =39-07% 

(2)  Ca  =  29.40 

(3)  K   =38.67 


O  =  58.49% 
S  =  23.56 
N=  13.88 


H=    2.44% 
O  =  47.04 
O  =  47.45 


MOLECULAR 
WEIGHT 

8l.O 
136.2 
1 61.2 


3.  The  molecular  weight  of  ammonia  is  17.06;  of  sulphur  dioxide 
is  64.06 ;  of  chlorine  is  70.9.    From  the  molecular  weight  calculate 
the  weight  of  i  1.  of  each  of  these  gases.    Compare  your  results  with 
the  table  on  the  back  cover  of  the  book. 

4.  From  the  molecular  weight  of  the  same  gases  calculate  the 
density  of  each,  referred  to  air  as  a  standard. 

5.  A  mixture  of  50  cc.  of  carbon  monoxide  and  50  cc.  of  oxygen 
was   exploded  in  a  eudiometer.     («)  What  gases  remained  in  the 
tube  after  the  explosion  ?    (£)  What  was  the  volume  of  each  ? 

6.  In  what  proportion  must  acetylene  and  oxygen  be  mixed  to 
produce  the  greatest  explosion  ? 

7.  Solve  Problem   18,  Chapter  XVII,   without   using   molecular 
weights.    Compare  your  results. 

8.  Solve  Problem    10,  Chapter  XVIII,  without  using  molecular 
weights.    Compare  your  results. 

9.  The  specific  heat  of  aluminium  is  0.214;  of  lead  is  0.031. 
From  these  specific  heats  calculate  the  atomic  weights  of  each  of 
the  elements. 


CHAPTER  XX 
THE  PHOSPHORUS  FAMILY 


SYMBOL 

ATOMIC 
WEIGHT 

DENSITY 

MELTING 
POINT 

Phosphorus   
Arsenic      ... 

P 

As 

3I.O 
7  C  o 

i.S 

57-5 

43-3° 

Antimony 

Sb 

/  J-w 
I  ''O  ^ 

•/  o 

6  7 

A1->° 

Bismuth     . 

Bi 

208.5 

u./ 
9.8 

4j- 
270° 

The  family.  The  elements  constituting  this  family 
belong  in  the  same  group  with  nitrogen  and  therefore 
resemble  it  in  a  general  way.  They  exhibit  a  regular  gra- 
dation of  physical  properties,  as  is  shown  in  the  above 
table.  The  same  general  gradation  is  also  found  in  their 
chemical  properties,  phosphorus  being  an  acid-forming 
element,  while  bismuth  is  essentially  a  metal.  The  other 
two  elements  are  intermediate  in  properties. 

Compounds.  In  general  the  elements  of  the  family  form 
compounds  having  similar  composition,  as  is  shown  in  the 
following  table  : 


PH3 
AsH3 
SbH3 

PC13 
AsCl3 
SbCl3 

PC15 
AsCl5 
SbCl5 

.  .  .  . 

BiCl3 

BiCl5 

As2Os 
Sb2(X 
Bi203 


AS2°5 

Sb205 
Bi205 


In  the  case  of  phosphorus,  arsenic,  and  antimony  the 
oxides  are  acid  anhydrides.  Salts  of  at  least  four  acids  of 
each  of  these  three  elements  are  known,  the  free  acid  in 

238 


THE  PHOSPHORUS  FAMILY  239 

some  instances  being  unstable.  The  relation  of  these  acids 
to  the  corresponding  anhydrides  may  be  illustrated  as  fol- 
lows, phosphorus  being  taken  as  an  example  : 

P2O3  +  3  H2O  =  2  H3PO3  (phosphorous  acid). 
P2O5  +  3  H2O  =  2  H3PO4  (phosphoric  acid). 
P2O5  +  2  H2O  =     H4P2O7  (pyrophosphoric  acid). 
P2O5  +     H2O  =  2  HPO3  (metaphosphoric  acid). 


PHOSPHORUS 

History.  The  element  phosphorus  was  discovered  by 
the  alchemist  Brand,  of  Hamburg,  in  1669,  while  searching 
for  the  philosopher's  stone.  Owing  to  its  peculiar  prop- 
erties and  the  secrecy  which  was  maintained  about  its 
preparation,  it  remained  a  very  rare  and  costly  substance 
until  the  demand  for  it  in  the  manufacture  of  matches 
brought  about  its  production  on  a  large  scale. 

Occurrence.  Owing  to  its  great  chemical  activity  phos- 
phorus never  occurs  free  in  nature.  In  the  form  of 
phosphates  it  is  very  abundant  and  widely  distributed. 
Phosphorite  and  sombrerite  are  mineral  forms  of  calcium 
phosphate,  while  apatite  consists  of  calcium  phosphate 
together  with  calcium  fluoride  or  chloride.  These  minerals 
form  very  large  deposits  and  are  extensively  mined  for 
use  as  fertilizers.  Calcium  phosphate  is-  a  constituent  of 
all  fertile  soil,  having  been  supplied  to  the  soil  by  the  dis- 
integration of  rocks  containing  it.  It  is  the  chief  mineral 
constituent  of  bones  of  animals,  and  bone  ash  is  therefore 
nearly  pure  calcium  phosphate. 

Preparation.  Phosphorus  is  now  manufactured  from  bone 
ash  or  a  pure  mineral  phosphate  by  heating  the  phosphate 
with  sand  and  carbon  in  an  electric  furnace.  The  materials 


240 


AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


are  fed  in  at  M  (Fig.  70)  by  the  feed  screw  F.    The  phos- 

phorus vapor  escapes  at  P 
and  is  condensed  under  water, 
while  the  calcium  silicate  is 
tapped  off  as  a  liquid  at  5. 
The  phosphorus  obtained  in 
this  way  is  quite  impure,  and 
is  purified  by  distillation. 

Explanation  of  the  reaction.  To 
understand  the  reaction  which  oc- 
curs, it  must  be  remembered  that  a 
volatile  acid  anhydride  is  expelled 
from  its  salts  when  heated  with  an 
anhydride  which  is  not  volatile. 

Thus,  when  sodium  carbonate  and  silicon  dioxide  are  heated  together 

the  following  reaction  takes  place  : 

Na2CO3  +  SiO2  =  Na2SiO3  +  CO2. 

Silicon  dioxide  is  a  less  volatile  anhydride  than  phosphoric  anhydride 
(P2O5),  and  when  strongly  heated  with  a  phosphate  the  phosphoric 
anhydride  is  driven  out,  thus  : 


FIG. 70 


Ca3(P04)2  +  3  Si02  =  3  CaSi0 


P2O5. 


If  carbon  is  added  before  the  heat  is  applied,  the  P2O5  is  reduced  to 
phosphorus  at  the  same  time,  according  to  the  equation 

P205  +  5  C  -  2  P  +  5  CO. 

Physical  properties.  The  purified  phosphorus  is  a  pale 
yellowish,  translucent,  waxy  solid  which  melts  at  43.3° 
and  boils  at  269°.  It  can  therefore  be  cast  into  any  con- 
venient form  under  warm  water,  and  is  usually  sold  in  the 
market  in  the  form  of  sticks.  It  is  quite  soft  and  can  be 
easily  cut  with  a  knife,  but  this  must  always  be  done  while 
the  element  is  covered  with  water,  since  it  is  extremely 
inflammable,  and  the  friction  of  the  knife  blade  is  almost 


THE  PHOSPHORUS  FAMILY  241 

sure  to  set  it  on  fire  if  cut  in  the  air.  It  is  not  soluble  in 
water,  but  is  freely  soluble  in  some  other  liquids,  notably 
in  carbon  disulphide.  Its  density  is  1.8. 

Chemical  properties.  Exposed  to  the  air  phosphorus 
slowly  combines  with  oxygen,  and  in  so  doing  emits  a  pale 
light,  or  phosphorescence,  which  can  be  seen  only  in  a  dark 
place.  The  heat  of  the  room  may  easily  raise  the  tempera- 
ture to  the  kindling  point  of  phosphorus,  when  it  burns  with 
a  sputtering  flame,  giving  off  dense  fumes  of  oxide  of  phos- 
phorus. It  burns  with  dazzling  brilliancy  in  oxygen,  and 
combines  directly  with  many  other  elements,  especially 
with  sulphur  and  the  halogens.  On  account  of  its  great 
affinity  for  oxygen  it  is  always  preserved  under  water. 

Phosphorus  is  very  poisonous,  from  0.2  to  0.3  gram  being 
a  fatal  dose.  Ground  up  with  flour  and  water  or  similar 
substances,  it  is  often  used  as  a  poison  for  rats  and  other 
vermin. 

Precaution.  The  heat  of  the  body  is  sufficient  to  raise  phosphorus 
above  its  kindling  temperature,  and  for  this  reason  it  should  always 
be  handled  with  forceps  and  never  with  the  bare  fingers.  .-Burns 
occasioned  by  it  are  very  painful  and  slow  in  healing. 

Red  phosphorus.  On  standing,  yellow  phosphorus  grad- 
ually undergoes  a  remarkable  change,  being  converted  into 
a  dark  red  powder  which  has  a  density  of  2.1.  It  no  longer 
takes  fire  easily,  neither  does  it  dissolve  in  carbon  disulphide. 
It  is  not  poisonous  and,  in  fact,  seems  to  be  an  entirely 
different  substance.  The  velocity  of  this  change  increases 
with  rise  in  temperature,  and  the  red  phosphorus  is  there- 
fore prepared  by  heating  the  yellow  just  below  the  boiling 
point  (25O°-3OO°).  When  distilled  and  quickly  condensed 
the  red  form  changes  back  to  the  yellow.  This  is  in  accord- 
ance with  the  general  rule  that  when  a  substance  capable 


242     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


of  existing  in  several  allotropic  forms  is  condensed  from  a 
gas  or  crystallized  from  the  liquid  state,  the  more  unstable 
variety  forms  first,  and  this  then  passes  into  the  more  stable 
forms. 

Matches.  The  chief  use  of  phosphorus  is  in  the  manufacture  of 
matches.  Common  matches  are  made  by  first  dipping  the  match  sticks 
into  some  inflammable  substance,  such  as  melted  paraffin,  and  after- 
ward into  a  paste  consisting  of  (i)  phosphorus,  (2)  some  oxidizing 
substance,  such  as  manganese  dioxide  or  potassium  chlorate,  and  (3) 
a  binding  material,  usually  some  kind  of  glue.  On  friction  the  phos- 
phorus is  ignited,  the  combustion  being  sustained  by  the  oxidizing 
agent  and  communicated  to  the  wood  by  the  burning  paraffin.  In 
sulphur  matches  the  paraffin  is  replaced  by  sulphur. 

In  safety  matches  red  phosphorus,  an  oxidizing  agent,  and  some 
gritty  material  such  as  emery  is  placed  on  the  side  of  the  box,  while 
the  match  tip  is  provided  as  before  with  an  oxidizing  agent  and  an 
easily  oxidized  substance,  usually  antimony  sulphide.  The  match 

cannot  be  ignited  easily  by  friction,  save  on  the 

prepared  surface. 

Compounds  of  phosphorus  with  hydro- 
gen.    Phosphorus    forms    several    com- 
pounds with  hydrogen,  the  best  known 
of  which  is  phosphine  (PH3)  analogous 
to  ammonia  (NH8). 

Preparation  of  phos- 
phine. Phosphine  is  usu- 
ally made  by  heating 
phosphorus  with  a  strong 
solution  of  potassium  hy- 
droxide, the  reaction  being 
a  complicated  one. 

The  experiment  can  be  conveniently  made  in  the  apparatus  shown 
in  Fig.  71.  A  strong  solution  of  potassium  hydroxide  together  with 
several  small  bits  of  phosphorus  are  placed  in  the  flask  A,  and  a 
current  of  coal  gas  is  passed  into  the  flask  through  the  tube  B  until 


FIG.  71 


THE  PHOSPHORUS  FAMILY 


243 


all  the  air  has  been  displaced.  The  gas  is  then  turned  off  and  the 
flask  is  heated.  Phosphine  is  formed  in  small  quantities  and  escapes 
through  the  delivery  tube,  the  exit  of  which  is  just  covered  by  the  water 
in  the  vessel  C.  Each  bubble  of  the  gas  as  it  escapes  into  the  air 
takes  fire,  and  the  product  of  combustion  (P2O5)  forms  beautiful  small 
rings,  which  float  unbroken  for  a  considerable  time  in  quiet  air.  The 
pure  phosphine  does  not  take  fire  spontaneously.  When  prepared  as 
directed  above,  impurities  are  present  which  impart  this  property. 

Properties.  Phosphine  is  a  gas  of  unpleasant  odor  and 
is  exceedingly  poisonous.  Like  ammonia  it  forms  salts 
with  the  halogen  acids.  Thus  we  have  phosphonium  chlo- 
ride (PH4C1)  analogous  to  ammonium  chloride  (NH4C1). 
The  phosphonium  salts  are  of  but  little  importance. 

Oxides  of  phosphorus.  Phosphorus  forms  two  well-known 
oxides,  —  the  trioxide  (P2O3)  and  the  pentoxide  (P2O5), 
sometimes  called  phosphoric  anhydride.  When  phosphorus 
burns  in  an  insufficient  supply  of  air  the  product  is  partially 
the  trioxide ;  in  oxygen  or  an  excess  of  air  the  pentoxide 
is  formed.  The  pentoxide  is  much  the  better  known  of  the 
two.  It  is  a  snow-white,  voluminous  powder  whose  most 
marked  property  is  its  great  attraction  for  water.  It  has 
no  chemical  action  upon  most  gases,  so  that  they  can  be 
very  thoroughly  dried  by  allowing  them  to  pass  through 
properly  arranged  vessels  containing  phosphorus  pentoxide. 

Acids  of  phosphorus.  The  important  acids  of  phosphorus 
are  the  following : 

H3PO3 phosphorous  acid. 

H3PO4 phosphoric  acid. 

H4P2O7 pyrophosphoric  acid. 

HPO3 metaphosphoric  acid. 

These  may  be  regarded  as  combinations  of  the  oxides  of 
phosphorus  with  water  according  to  the  equations  given  in 
the  discussion  of  the  characteristics  of  the  family. 


244     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

1.  Phosphorous  acid  (H3PO3).    Neither  the  acid  nor  its 
salts  are  at  all  frequently  met  with  in  chemical  operations. 
It  can  be  easily  obtained,  however,  in  the  form  of  transpar- 
ent crystals  when  phosphorus  trichloride  is  treated  with 
water  and  the  resulting  solution  is  evaporated  : 

PC13  +  3  H2O  =  H3PO3  +  3  HC1. 

Its  most  interesting  property  is  its  tendency  to  take  up 
oxygen  and  pass  over  into  phosphoric  acid. 

2.  Orthophosphoric    acid    (phosphoric    acid]     (H3PO4). 
This  acid  can  be  obtained  by  dissolving  phosphorus  pen- 
toxide  in  boiling  water,  as  represented  in  the  equation 

P205  +  3  H20  =  2  H3P04. 

It  is  usually  made  by  treating  calcium  phosphate  with  con- 
centrated sulphuric  acid.  The  calcium  sulphate  produced 
in  the  reaction  is  nearly  insoluble,  and  can  be  filtered  off, 
leaving  the  phosphoric  acid  in  solution.  Very  pure  acid  is 
made  by  oxidizing  phosphorus  with  nitric  acid.  It  forms 
large  colorless  crystals  which  are  exceedingly  soluble  in 
water.  Being  a  tribasic  acid,  it  forms  acid  as  well  as 
normal  salts.  Thus  the  following  compounds  of  sodium 
are  known : 

NaH2PO4 monosodium  hydrogen  phosphate. 

Na2HPO4 disodium  hydrogen  phosphate. 

Na3PO4 normal  sodium  phosphate. 

These  salts  are  sometimes  called  respectively  primary, 
secondary,  and  tertiary  phosphates.  They  may  be  pre- 
pared by  bringing  together  phosphoric  acid  and  appro- 
priate quantities  of  sodium  hydroxide.  Phosphoric  acid 
also  forms  mixed  salts,  that  is,  salts  containing  two  differ- 
ent metals.  The  most  familiar  compound  of  this  kind  is 
microcosmic  salt,  which  has  the  formula  Na(NH4)HPO4. 


THE  PHOSPHORUS  FAMILY 


245 


Orthophosphates.  The  orthophosphates  form  an  impor- 
tant class  of  salts.  The  normal  salts  are  nearly  all  insoluble 
and  many  of  them  occur  in  nature.  The  secondary  phos- 
phates are  as  a  rule  insoluble,  while  most  of  the  primary 
salts  are  soluble. 

3.  Pyrophosphoric  acid  (H4P2O7).     On  heating  orthophosphoric 
acid  to  about  225°  pyrophosphoric  acid  is  formed  in  accordance  with 
the  following  equation : 

2  H3P04  -  H4P207  +  H20. 

It  is  a  white  crystalline  solid.  Its  salts  can  be  prepared  by  heating 
a  secondary  phosphate : 

2  Na2HPO4  =  Na4P2O7  +  H2O. 

4.  Metaphosphoric    acid    {glacial    phosphoric     acid}     (HPO8). 
This  acid  is  formed  when  orthophosphoric  acid  is  heated  above  400°  : 

H3PO4  =  HPO3  +  H2O. 

It  is  also  formed  when  phosphorus  pentoxide  is  treated  with  cold 
water :  P2O5  +  H20  -  2  HPO3. 

It  is  a  white  crystalline  solid,  and  is  so  stable  towards  heat  that  it  can 
be  fused  and  even  volatilized  without  decomposition.  On  cooling 
from  the  fused  state  it  forms  a  glassy  solid,  and  on  this  account  is 
often  called  glacial  phosphoric  acid.  It  possesses  the  property  of 
dissolving  small  quantities  of  metallic  oxides,  with  the  formation  of 
compounds  which,  in  the  case  of  certain  metals,  have  characteristic 
colors.  It  is  therefore  used  in  the  detection  of  these  metals. 

While  the  secondary  phosphates,  on  heating,  give  salts  of  pyro- 
phosphoric acid,  the  primary  phosphates  yield  salts  of  metaphos- 
phoric  acid.  The  equations  representing  these  reactions  are  as 
follows :  2  Na2Hp04  =  Na4P207  +  H2O, 

NaH2PO4  =  NaPO3  +  H2O. 

Fertilizers.  When  crops  are  produced  year  after  year 
on  the  same  field  certain  constituents  of  the  soil  essential 
to  plant  growth  are  removed,  and  the  soil  becomes  impov- 
erished and  unproductive.  To  make  the  land  once  more 


246     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

fertile  these  constituents  must  be  replaced.  The  calcium 
phosphate  of  the  mineral  deposits  or  of  bone  ash  serves 
well  as  a  material  for  restoring  phosphorus  to  soils  exhausted 
of  that  essential  element ;  but  a  more  soluble  substance, 
which  the  plants  can  more  readily  assimilate,  is  desirable. 
It  is  better,  therefore,  to  convert  the  insoluble  calcium  phos- 
phate into  the  soluble  primary  phosphate  before  it  is  applied 
as  fertilizer.  It  will  be  seen  by  reference  to  the  formulas 
for  the  orthophosphates  (see  page  244)  that  in  a  pri- 
mary phosphate  only  one  hydrogen  atom  of  phosphoric  acid 
is  replaced  by  a  metal.  Since  the  calcium  atom  always 
replaces  two  hydrogen  atoms,  it  might  be  thought  that 
there  could  be  no  primary  calcium  phosphate ;  but  if  the 
calcium  atom  replaces  one  hydrogen  atom  from  each  of 
two  molecules  of  phosphoric  acid,  the  salt  Ca(H2PO4)2 
will  result,  and  this  is  a  primary  phosphate.  It  can  be 
made  by  treatment  of  the  normal  phosphate  with  the  nec- 
essary amount  of  sulphuric  acid,  calcium  sulphate  being 
formed  at  the  same  time,  thus  : 

Ca3(PO4)2  +  2  H2SO4  =  Ca(H2PO4)2  +  2  CaSO4. 

The  resulting  mixture  is  a  powder,  which  is  sold  as  a  fer- 
tilizer under  the  name  of  "  superphosphate  of  lime." 

ARSENIC 

Occurrence.  Arsenic  occurs  in  considerable  quantities 
in  nature  as  the  native  element,  as  the  sulphides  realgar 
(As2S2)  and  orpiment  (As2S3),  as  oxide  (As2O3),  and  as  a 
constituent  of  many  metallic  sulphides,  such  as  arsenopyrite 
(FeAsS). 

Preparation.  The  element  is  prepared  by  purifying  the 
native  arsenic,  or  by  heating  the  arsenopyrite  in  iron  tubes, 


THE  PHOSPHORUS  FAMILY  247 

out  of  contact  with  air,  when  the  reaction  expressed  by  the 
following  equation  occurs  : 

FeAsS  -  FeS  +  As. 

The  arsenic,  being  volatile,  condenses  in  chambers  con- 
nected with  the  heated  tubes.  It  is  also  made  from  the 
oxide  by  reduction  with  carbon : 

2  As2O3  +  3C=4As  +  3  CO2. 

Properties.  Arsenic  is  a  steel-gray,  metallic-looking  sub- 
tance  of  density  5.73.  Though  resembling  metals  in 
appearance,  it  is  quite  brittle,  being  easily  powdered  in  a 
mortar.  When  strongly  heated  it  sublimes,  that  is,  it 
passes  into  a  vapor  without  melting,  and  condenses  again 
to  a  crystalline  solid  when  the  vapor  is  cooled.  Like  phos- 
phorus it  can  be  obtained  in  several  allotropic  forms.  It 
alloys  readily  with  some  of  the  metals,  and  finds  its  chief 
use  as  an  alloy  with  lead,  which  is  used  for  making  shot, 
the  alloy  being  harder  than  pure  lead.  When  heated  on 
charcoal  with  the  blowpipe  it  is  converted  into  an  oxide 
which  volatilizes,  leaving  the  charcoal  unstained  by  any 
oxide  coating.  It  burns  readily  in  chlorine  gas,  forming 
arsenic  trichloride,  — 

As  +  3  Cl  -  AsCl3. 

Unlike  most  of  its  compounds,  the  element  itself  is  not 
poisonous. 

Arsine  (AsH3).  When  any  compound  containing  arsenic 
is  brought  into  the  presence  of  nascent  hydrogen,  arsine 
(AsH3),  corresponding  to  phosphine  and  ammonia,  is 
formed.  The  reaction  when  oxide  of  arsenic  is  so  treated 

As2O3  +  12  H  =  2  AsH3  +  3  H2O. 


248     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


Arsine  is  a  gas  with  a  peculiar  garlic-like  odor,  and  is 
intensely  poisonous.  A  single  bubble  of  pure  gas  has  been 
known  to  prove  fatal.  It  is  an  unstable  compound,  de- 
composing into  its  elements  when  heated  to  a  moderate 
temperature.  It  is  combustible,  burning  with  a  pale  bluish- 
white  flame  to  form  arsenic  trioxide  and  water  when  air 

is  in  excess:       .    TT  TT  ~. 

2  AsH3  +  6  O=  As2O3  +  3  H2O. 

When  the  supply  of  air  is  deficient  water  and  metallic 
arsenic  are  formed  : 

2  AsH3  +30  =  3  H2O  +  2  As. 

These  reactions  make  the  detection  of  even  minute  quanti- 
ties of  arsenic  a  very  easy  problem. 

Marsh's  test  for  arsenic.  The  method  devised  by  Marsh  for 
detecting  arsenic  is  most  frequently  used,  the  apparatus  being  shown 

in  Fig.  72.  Hydrogen  is  generated 
in  the  flask  A  by  the  action  of  dilute 
sulphuric  acid  on  zinc,  is  dried  by 
passing  over  calcium  chloride  in  the 
tube  B,  and  after  passing  through 


o 


FIG.  72 


the  hard-glass  tube  C  is  ignited  at  the  jet  D.  If  a  substance  contain- 
ing arsenic  is  now  introduced  into  the  generator  A,  the  arsenic  is 
converted  into  arsine  by  the  action  of  the  nascent  hydrogen,  and 


THE  PHOSPHORUS  FAMILY  249 

passes  to  the  jet  along  with  the  hydrogen.  If  the  tube  C  is  strongly 
heated  at  some  point  near  the  middle,  the  arsine  is  decomposed  while 
passing  this  point  and  the  arsenic  is  deposited  just  beyond  the  heated 
point  in  the  form  of  a  shining,  brownish-black  mirror.  If  the  tube  is 
not  heated,  the  arsine  burns  along  with  the  hydrogen  at  the  jet.  Under 
these  conditions  a  small  porcelain  dish  crowded  down  into  the  flame  is 
blackened  by  a  spot  of  metallic  arsenic,  for  the  arsine  is  decomposed 
by  the  heat  of  the  flame,  and  the  arsenic,  cooled  below  its  kindling 
temperature  by  the  cold  porcelain,  deposits  upon  it  as  a  black  spot. 
Antimony  conducts  itself  in  the  same  way  as  arsenic,  but  the  antimony 
deposit  is  more  sooty  in  appearance.  The  two  can  also  be  distin- 
guished by  the  fact  that  sodium  hypochlorite  (NaCIO)  dissolves  the 
arsenic  deposit,  but  not  that  formed  by  antimony. 

Oxides  of  arsenic.  Arsenic  forms  two  oxides,  As2O3  and 
As2O5,  corresponding  to  those  of  phosphorus.  Of  these 
arsenious  oxide,  or  arsenic  trioxide  (As2O3),  is  much  better 
known,  and  is  the  substance  usually  called  white  arsenic, 
or  merely  arsenic.  It  is  found  as  a  mineral,  but  is  usually 
obtained  as  a  by-product  in  burning  pyrite  in  the  sulphuric- 
acid  industry.  The  pyrite  has  a  small  amount  of  arseno- 
pyrite  in  it,  and  when  this  is  burned  arsenious  oxide  is 
formed  as  a  vapor  together  with  sulphur  dioxide  : 

2  FeAsS  +  10  O  =  Fe2O3  +  As2O3  +  2  SO2. 

The  arsenious  oxide  is  condensed  in  appropriate  chambers. 
It  is  a  rather  heavy  substance,  obtained  either  as  a  crystal- 
line powder  or  as  large,  vitreous  lumps,  'resembling  lumps 
of  porcelain  in  appearance.  It  is  very  poisonous,  from  0.2 
to  o.  3  g.  being  a  fatal  dose.  It  is  frequently  given  as  a  poison, 
since  it  is  nearly  tasteless  and  does  not  act  very  rapidly. 
This  slow  action  is  due  to  the  fact  that  it  is  not  very 
soluble,  and  hence  is  absorbed  slowly  by  the  system. 
Arsenious  oxide  is  also  used  as  a  chemical  reagent  in. 
glass  making  and  in  the  dye  industry. 


250     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Acids  of  arsenic.  Like  the  corresponding  oxides  of  phosphorus, 
the  oxides  of  arsenic  are  acid  anhydrides.  In  solution  they  combine 
with  bases  to  form  salts,  corresponding  to  the  salts  of  the  acids  of 
phosphorus.  Thus  we  have  salts  of  the  following  acids  : 

H3AsO3    .......  arsenious  acid. 

H3AsO4    .     .  '  !     .     .     .     .  orthoarsenic  acid. 

H4As2O7  .......  pyroarsenic  acid. 

HAsO,  metarsenic  acid. 

o 

Several  other  acids  of  arsenic  are  also  known.  Not  all  of  these 
can  be  obtained  as  free  acids,  since  they  tend  to  lose  water  and  form 
the  oxides.  Thus,  instead  of  obtaining  arsenious  acid  (H3AsO3),  the 
oxide  As2O3  is  obtained  : 

2  H3AsO3  =  As2O3  +  3  H2O. 

Salts  of  all  the  acids  are  known,  however,  and  some  of  them 
have  commercial  value.  Most  of  them  are  insoluble,  and  some  of  the 
copper  salts,  which  are  green,  are  used  as  pigments.  Paris  green, 
which  has  a  complicated  formula,  is  a  well-known  insecticide. 

Antidote  for  arsenical  poisoning.  The  most  efficient  antidote  for 
arsenic  poisoning  is  ferric  hydroxide.  It  is  prepared  as  needed, 
according  to  the  equation 

Fe2  (S04)3  +  3  Mg  (OH)2  =  2  Fe  (OH)3  +  3  MgSO4. 

Sulphides  of  arsenic.  When  hydrogen  sulphide  is  passed  into  an 
acidified  solution  containing  an  arsenic  compound  the  arsenic  is 
precipitated  as  a  bright  yellow  sulphide,  thus  : 

2  H3AsO3  +  3  H2S  =  As2S3  +6  H2O, 
2  H3AsO4  +  5  H2S  =  As2S5  +  8  H2O. 

In  this  respect  arsenic  resembles  the  metallic  elements,  many  of 
which  produce  sulphides  under  similar  conditions.  The  sulphides  of 
arsenic,  both  those  produced  artificially  and  those  found  in  nature, 
are  used  as  yellow  pigments. 

ANTIMONY 

Occurrence.  Antimony  occurs  in  nature  chiefly  as  the 
sulphide  (Sb2S3),  called  stibnite,  though  it  is  also  found 
as  oxide  and  as  a  constituent  of  many  complex  minerals. 


THE  PHOSPHORUS  FAMILY 


251 


Preparation.  Antimony  is  prepared  from  the  sulphide 
in  a  very  simple  manner.  The  sulphide  is  melted  with 
scrap  iron  in  a  furnace,  when  the  iron  combines  with  the 
sulphur  to  form  a  slag,  or  liquid  layer  of  melted  iron  sul- 
phide, while  the  heavier  liquid,  antimony,  settles  to  the 
bottom  and  is  drawn  off  from  time  to  time.  The  reaction 
involved  is  represented  by  the  equation 

Sb2S3  +  3  Fe  =  2  Sb  +  3  FeS. 

Physical  properties.  Antimony  is  a  bluish-white,  metallic- 
looking  substance  whose  density  is  6.7.  It  is  highly  crys- 
talline, hard,  and  very  brittle.  It  has  a  rather  low  melting 
point  (432°)  and  expands  very  noticeably  on  solidifying. 

Chemical  properties.  In  chemical  properties  antimony 
resembles  arsenic  in  many  particulars.  It  forms  the  oxides 
Sb2O3  and  Sb2O6,  and  in  addition  Sb2O4.  It  combines 
with  the  halogen  elements  with  great  energy,  burning 
brilliantly  in  chlorine  to  form  antimony  trichloride  (SbCl3). 
When  heated  on  charcoal  with  the  blowpipe  it  is  oxidized 
and  forms  a  coating  of  antimony  oxide  on  the  charcoal 
which  has  a  characteristic  bluish-white  color. 

Stibine  (SbH3).  The  gas  stibine  (SbH3)  is  formed  under 
conditions  which  are  very  similar  to  those  which  produce 
arsine,  and  it  closely  resembles  the  latter  compound,  though 
it  is  still  less  stable.  It  is  very  poisonous. 

Acids  of  antimony.  The  oxides  Sb2O3  and  Sb2O5  are  weak  acid 
anhydrides  and  are  capable  of  forming  two  series  of  acids  corre- 
sponding in  formulas  to  the  acids  of  phosphorus  and  arsenic.  They 
are  much  weaker,  however,  and  are  of  little  practical  importance. 

Sulphides  of  antimony.    Antimony  resembles  arsenic  in  that  hydro- 
gen sulphide  precipitates  it  as  a  sulphide  when  conducted  into  an 
acidified  solution  containing  an  antimony  compound : 
2  SbCl3  +  3  H2S  =  Sb2S3  +  6  HC1, 
2  SbCl5  +  5  H2S  =  Sb2S5  +  10  HC1. 


252     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

The  two  sulphides  of  antimony  are  called  the  trisulphide  and  the 
pentasulphide  respectively.  When  prepared  in  this  way  they  are 
orange-colored  substances,  though  the  mineral  stibnite  is  black. 

Metallic  properties  of  antimony.  The  physical  properties 
of  the  element  are  those  of  a  metal,  and  the  fact  that  its 
sulphide  is  precipitated  by  hydrogen  sulphide  shows  that  it 
acts  like  a  metal  in  a  chemical  way.  Many  other  reactions 
show  that  antimony  has  more  of  the  properties  of  a  metal 
than  of  a  non-metal.  The  compound  Sb(OH)3,  correspond- 
ing to  arsenious  acid,  while  able  to  act  as  a  weak  acid  is 
also  able  to  act  as  a  weak  base  with  strong  acids.  For 
example,  when  treated  with  concentrated  hydrochloric  acid 
antimony  chloride  is  formed  : 

Sb(OH)3  +  3  HC1  =  SbCl3  +  3  H2O. 

A  number  of  elements  act  in  this  same  way,  their  hy- 
droxides under  some  conditions  being  weak  acids  and 
under  others  weak  bases. 

ALLOYS 

Some  metals  when  melted  together  thoroughly  intermix, 
and  on  cooling  form  a  homogeneous,  metallic-appearing 
substance  called  an  alloy.  Not  all  metals  will  mix  in  this 
way,  and  in  some  cases  definite  chemical  compounds  are 
formed  and  separate  out  as  the  mixture  solidifies,  thus 
destroying  the  uniform  quality  of  the  alloy.  In  general  the 
melting  point  of  the  alloy  is  below  the  average  of  the 
melting  points  of  its  constituents,  and  it  is  often  lower 
than  any  one  of  them. 

Antimony  forms  alloys  with  many  of  the  metals,  and  its 
chief  commercial  use  is  for  such  purposes.  It  imparts  to 
its  alloys  high  density,  rather  low  melting  point,  and  the 


THE  PHOSPHORUS  FAMILY  253 

property  of  expanding  on  solidification.  Such  an  alloy  is 
especially  useful  in  type  founding,  where  fine  lines  are  to 
be  reproduced  on  a  cast.  Type  metal  consists  of  antimony, 
lead,  and  tin.  Babbitt  metal,  used  for  journal  bearings  in 
machinery,  contains  the  same  metals  in  a  different  pro- 
portion together  with  a  small  percentage  of  copper. 

BISMUTH 

Occurrence.  Bismuth  is  usually  found  in  the  uncombined 
form  in  nature.  It  also  occurs  as  oxide  and  sulphide.  Most 
of  the  bismuth  of  commerce  comes  from  Saxony,  and  from 
Mexico  and  Colorado,  but  it  is  not  an  abundant  element. 

Preparation.  It  is  prepared  by  merely  heating  the  ore 
containing  the  native  bismuth  and  allowing  the  melted 
metal  to  run  out  into  suitable  vessels.  Other  ores  are  con- 
verted into  oxides  and  reduced  by  heating  with  carbon. 

Physical  properties.  Bismuth  is  a  heavy,  crystalline, 
brittle  metal  nearly  the  color  of  silver,  but  with  a  slightly 
rosy  tint  which  distinguishes  it  from  other  metals.  It  melts 
at  a  low  temperature  (270°)  and  has  a  density  of  9.8.  It 
is  not  acted  upon  by  the  air  at  ordinary  temperatures. 

Chemical  properties.  When  heated  with  the  blowpipe  on 
charcoal,  bismuth  gives  a  coating  of  the  oxide  Bi2O3.  This 
has  a  yellowish-brown  color  which  easily  distinguishes  it 
from  the  oxides  formed  by  other  metals.  It  combines  very 
readily  with  the  halogen  elements,  powdered  bismuth  burn- 
ing readily  in  chlorine.  It  is  not  very  easily  acted  upon  by 
hydrochloric  acid,  but  nitric  and  sulphuric  acids  act  upon 
it  in  the  same  way  that  they  do  upon  copper. 

Uses.  Bismuth  finds  its  chief  use  as  a  constituent  of 
alloys,  particularly  in  those  of  low  melting  point.  Some 


254     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

of  these  melt  in  hot  water.  For  example,  Wood's  metal, 
consisting  of  bismuth,  lead,  tin,  and  cadmium,  melts  at 
60.5°. 

Compounds  of  bismuth.  Unlike  the  other  elements  of 
this  group,  bismuth  has  almost  no  acid  properties.  Its 
chief  oxide,  Bi2O3,  is  basic  in  its  properties.  It  dissolves 
in  strong  acids  and  forms  salts  of  bismuth  : 

Bi2O3  +  6  HC1  =  2  BiCl3  +  3  H2O, 

Bi2O3  +  6  HNO3  =  2  Bi(NO3)3  +  3  H2O. 

The  nitrate  and  chloride  of  bismuth  can  be  obtained  as 
well-formed  colorless  crystals.  When  treated  with  water 
the  salts  are  decomposed  in  the  manner  explained  in  the 
following  paragraph. 

HYDROLYSIS 

Many  salts  such  as  those  of  antimony  and  bismuth  form 
solutions  which  are  somewhat  acid  in  reaction,  and  must 
therefore  contain  hydrogen  ions.  This  is  accounted  for 
by  the  same  principle  suggested  to  explain  the  fact  that 
solutions  of  potassium  cyanide  are  alkaline  in  reaction 
(p.  210).  Water  forms  an  appreciable  number  of  hydrogen 
and  hydroxyl  ions,  and  very  weak  bases  such  as  bismuth 
hydroxide  are  dissociated  to  but  a  very  slight  extent.  When 
Bi+++  ions  from  bismuth  chloride,  which  dissociates  very 
readily,  are  brought  in  contact  with  the  OH~  ions  from 
water,  the  two  come  to  the  equilibrium  expressed  in  the 
equation  Bi+++  +  3  QH-  _^  Bi(OH)3. 

For  every  hydroxyl  ion  removed  from  the  solution  in  this 
way  a  hydrogen  ion  is  left  free,  and  the  solution  becomes 
acid  in  reaction. 


THE  PHOSPHORUS  FAMILY  255 

Reactions  of  this  kind  and  that  described  under  potas- 
sium cyanide  are  called  hydrolysis. 

DEFINITION  :  Hydrolysis  is  the  action  of  water  upon  a 
salt  to  form  an  acid  and  a  base,  one  of  which  is  very 
slightly  dissociated. 

Conditions  favoring  hydrolysis.  While  hydrolysis  is  primarily  due 
to  the  slight  extent  to  which  either  the  acid  or  the  base  formed  is 
dissociated,  several  other  factors  have  an  influence  upon  the  extent 
to  which  it  will  take  place. 

1 .  Influence  of  mass.    Since  hydrolysis  is  a  reversible  reaction, 
the  relative  masses  of  the  reacting  substances  influence  the  point  at 
which  equilibrium  will  be  reached.     In  the  equilibrium 

BiCl3  +  3  H20  ^  Bi(OH)3  +  3  HC1 

the  addition  of  more  water  will  result  in  the  formation  of  more 
bismuth  hydroxide  and  hydrochloric  acid.  The  addition  of  more 
hydrochloric  acid  will  convert  some  of  the  bismuth  hydroxide  into 
bismuth  chloride. 

2.  Formation  of  insoluble  substances.    When  one  of  the  products 
of  hydrolysis  is  nearly  insoluble  in  water  the  solution  will  become 
saturated  with  it  as  soon  as  a  very  little  has  been  formed.    All  in 
excess  of  this  will  precipitate,  and  the  reaction  will  go  on  until  the 
acid  set  free  increases  sufficiently  to  bring  about  an  equilibrium. 
Thus  a  considerable  amount  of  bismuth  and  antimony  hydroxides  are 
precipitated  when  water  is  added  to  the  chlorides  of  these  elements. 
The   greater   the   dilution   the    more   hydroxide    precipitates.     The 
addition  of  hydrochloric  acid  in  considerable  quantity  will,  however, 
redissolve  the  precipitate. 

Partial  hydrolysis.  In  many  cases  the  hydrolysis  of  a  salt  is  only 
partial,  resulting  in  the  formation  of  basic  sa'lts  instead  of  the  free 
base.  Most  of  these  basic  salts  are  insoluble  in  water,  which 
accounts  for  their  ready  formation.  Thus  bismuth  chloride  may 
hydrolyze  by  successive  steps,  as  shown  in  the  equations 

BiCl3  +     H2O  =  Bi(OH)Cl2  +  HC1, 
BiCl3  +  2  H2O  =  Bi(OH)2Cl  +  2  HC1, 
BiCl3  +  3  H20  =  Bi(OH)3  +  3  HC1. 

The  basic  salt  so  formed  may  also  lose  water,  as  shown   in  the 

equation  Bi(OH)2Cl  =  BiOCl  +  H2O. 


256     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

The  salt  represented  in  the  last  equation  is  sometimes  called  bis- 
muth oxychloride,  or  bismuthyl  chloride.  The  corresponding  nitrate, 
BiONO3,  is  largely  used  in  medicine  under  the  name  of  subnitrate 
of  bismuth.  In  these  two  compounds  the  group  of  atoms,  BiO,  acts 
as  a  univalent  metallic  radical  and  is  called  bismuthyL  Similar  basic 
salts  are  formed  by  the  hydrolysis  of  antimony  salts. 


EXERCISES 

1.  Name  all  the  elements  so  far  studied  which  possess  allotropic 
forms. 

2.  What  compounds  would  you  expect  phosphorus  to  form  with 
bromine  and  iodine  ?    Write   the   equations  showing  the  action  of 
water  on  these  compounds. 

3.  In  the  preparation  of  phosphine,  why  is  coal  gas  passed  into 
the  flask?    What  other  gases  would  serve  the  same  purpose? 

4.  Give  the  formula  for  the  salt  which  phosphine  forms  with 
hydriodic  acid.     Give  the  name  of  the  compound. 

5.  Could  phosphoric  acid  be  substituted  for  sulphuric  acid  in 
the  preparation  of  the  common  acids  ? 

6.  Write  the  equations  for  the  preparation  of  the  three  sodium 
salts  of  orthophosphoric  acid. 

7.  Why  does  a  solution  of  disodium  hydrogen  phosphate  react 
alkaline? 

8.  On  the  supposition  that  bone  ash  is  pure  calcium  phosphate, 
what  weight  of  it  would  be  required  in  the  preparation  of  I  kg.  of 
phosphorus  ? 

9.  If  arsenopyrite  is  heated  in  a  current  of  air,  what  products 
are  formed  ? 

10.  («)  Write  equations  for  the  complete  combustion  of  hydro- 
sulphuric  acid,  methane,  and  arsine.  (£)  In  what  respects  are  the 
reactions  similar  ? 

•11.  Write  the  equations  for  all  the  reactions  involved  in  Marsh's 
test  for  arsenic. 

12.  Write  the  names  and  formulas  for  the  acids  of  antimony. 

13.  Write  the  equations  showing  the  hydrolysis  of  antimony  tri- 
chloride ;  of  bismuth  nitrate. 

14.  In  what  respects  does  nitrogen  resemble  the  members  of  the 
phosphorus  family  ? 


CHAPTER    XXI 
SILICON,  TITANIUM,  BORON 


SYMBOL 

ATOMIC 
WEIGHT 

DENSITY 

CHLORIDES 

OXIDES 

Silicon     »•";-'-. 

Si 

28  4 

2.1  C 

SiCl4 

SiO2 

Titanium 

Ti 

48  I 

-?  c 

TiCl4 

TiO2 

Boron     .    ...    .... 

B 

II.O 

2-45 

BC13 

B203 

General.  Each  of  the  three  elements,  silicon,  titanium, 
and  boron,  belongs  to  a  separate  periodic  family,  but  they 
occur  near  together  in  the  periodic  grouping  and  are  very 
similar  in  both  physical  and  chemical  properties.  Since 
the  other  elements  in  their  families  are  either  so  rare  that 
they  cannot  be  studied  in  detail,  or  are  best  understood  in 
connection,  with  other  elements,  it  is  convenient  to  con- 
sider these  three  together  at  this  point. 

The  three  elements  are  very  difficult  to  obtain  in  the  free 
state,  owing  to  their  strong  attraction  for  other  elements. 
They  can  be  prepared  by  the  action  of 'aluminium  or  mag- 
nesium on  their  oxides  and  in  impure  state  by  reduction 
with  carbon  in  an  electric  furnace.  They  are  very  hard 
and  melt  only  at  the  highest  temperatures.  At  ordinary 
temperatures  they  are  not  attacked  by  oxygen,  but  when 
strongly  heated  they  burn  with  great  brilliancy.  Silicon 
and  boron  are  not  attacked  by  acids  under  ordinary  con- 
ditions ;  titanium  is  easily  dissolved  by  them. 

257 


258     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

SILICON 

Occurrence.  Next  to  oxygen  silicon  is  the  most  abun- 
dant element.  It  does  not  occur  free  in  nature,  but  its 
compounds  are  very  abundant  and  of  the  greatest  impor- 
tance. It  occurs  almost  entirely  in  combination  with  oxygen 
as  silicon  dioxide  (SiO2),  often  called  silica,  or  with  oxygen 
and  various  metals  in  the  form  of  salts  of  silicic  acids,  or 
silicates.  These  compounds  form  a  large  fraction  of  the 
earth's  crust.  Most  plants  absorb  small  amounts  of  silica 
from  the  soil,  and  it  is  also  found  in  minute  quantities  in 
animal  organisms. 

Preparation.  The  element  is  most  easily  prepared  by 
reducing  pure  powdered  quartz  with  magnesium  powder : 

SiO2  +  2  Mg  =  2  MgO  +  Si. 

Properties.  As  would  be  expected  from  its  place  in  the 
periodic  table,  silicon  resembles  carbon  in  many  respects. 
It  can  be  obtained  in  several  allotropic  forms,  correspond- 
ing to  those  of  carbon.  The  crystallized  form  is  very  hard, 
and  is  inactive  toward  reagents.  The  amorphous  variety 
has,  in  general,  properties  more  similar  to  charcoal. 

Compounds  of  silicon  with  hydrogen  and  the  halogens. 
Silicon  hydride  (SiH4)  corresponds  in  formula  to  methane 
(CH4),  but  its  properties  are  more  like  those  of  phosphine 
(PH3).  It  is  a  very  inflammable  gas  of  disagreeable  odor, 
and,  as  ordinarily  prepared,  takes  fire  spontaneously"  on 
account  of  the  presence  of  impurities. 

Silicon  combines  with  the  elements  of  the  chlorine  family 
to  form  such  compounds  as  SiCl4  and  SiF4.  Of  these 
silicon  fluoride  is  the  most  familiar  and  interesting.  As 
stated  in  the  discussion  of  fluorine,  it  is  formed  when 


SILICON,  TITANIUM,  BORON  259 

hydrofluoric  acid  acts  upon  silicon  dioxide  or  a  silicate. 
With  silica  the  reaction  is  thus  expressed  : 

SiO2  +  4  HF  =  SiF4  +  2  H2O. 

It  is  a  very  volatile,  invisible,  poisonous  gas.  In  contact 
with  water  it  is  partially  decomposed,  as  shown  in  the 
equation 

SiF4  +  4  H2O  =  4  HF  +  Si(OH)4. 

The  hydrofluoric  acid  so  formed  combines  with  an  addi- 
tional amount  of  silicon  fluoride,  forming  the  complex 
fluosilicic  acid  (H2SiF6),  thus  : 

2  HF  +  SiF4  -  H2SiF6. 

Silicides.  As  the  name  indicates,  silicides  are  binary 
compounds  consisting  of  silicon  and  some  other  element. 
They  are  very  stable  at  high  temperatures,  and  are  usually 
made  by  heating  the  appropriate  substances  in  an  electric 
furnace.  The  most  important  one  is  carborundum,  which  is 
a  silicide  of  carbon  of  the  formula  CSi.  It  is  made  by  heat- 
ing coke  and  sand,  which  is  a  form  of  silicon  dioxide,  in  an 
electric  furnace,  the  process  being  extensively  carried  on 
at  Niagara  Falls.  The  following  equation  represents  the 

reaction 

SiO2  +  3  C  =  CSi  +  2  CO. 

The  substance  so  prepared  consists  of  beautiful  purplish- 
black  crystals,  which  are  very  hard.  Carborundum  is  used 
as  an  abrasive,  that  is,  as  a  material  for  grinding  and  polish- 
ing very  hard  substances.  Ferrosilicon  is  a  silicide  of  iron 
alloyed  with  an  excess  of  iron,  which  finds  extensive  use  in 
the  manufacture  of  certain  kinds  of  steel. 


260     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Manufacture  of  carborundum.  The  mixture  of  materials  is  heated  in 
a  large  resistance  furnace  for  about  thirty-six  hours.  After  the  reaction 

is  completed  there  is 
left  a  core  of  graph- 
ite G.  Surrounding 
this  core  is  a  layer 
of  crystallized  car- 
borundum C,  about 
1 6  in.  thick.  Out- 

FlG  side  this  is   a   shell 

of  amorphous  car- 
borundum A.  The  remaining  materials  M  are  unchanged  and 
are  used  for  a  new  charge. 

Silicon  dioxide  (silica]  (SiO2).  This  substance  is  found 
in  a  great  variety  of  forms  in  nature,  both  in  the  amor- 
phous and  in  the  crystalline  condition.  In  the  form  of  quartz 
it  is  found  in  beautifully  formed  six-sided  prisms,  sometimes 
of  great  size.  When  pure  it  is  perfectly  transparent  and 
colorless.  Some  colored  varieties  are  given  special  names, 
as  amethyst  (violet),  rose  quartz  (pale  pink),  smoky  or 
milky  quartz  (colored  and  opaque).  Other  varieties  of 
silicon  dioxide,  some  of  which  also  contain  water,  are 
chalcedony,  onyx,  jasper,  opal,  agate,  and  flint.  Sand  and 
sandstone  are  largely  silicon  dioxide. 

Properties.  As  obtained  by  chemical  processes  silicon 
dioxide  is  an  amorphous  white  powder.  In  the  crystallized 
state  it  is  very  hard  and  has  a  density  of  2.6.  It  is 
insoluble  in  water  and  in  most  chemical  reagents,  and  re- 
quires the  hottest  oxyhydrogen  flame  for  fusion.  Acids, 
excepting  hydrofluoric  acid,  have  little  action  on  it,  and  it 
requires  the  most  energetic  reducing  agents  to  deprive  it 
of  oxygen.  It  is  the  anhydride  of  an  acid,  and  consequently 
it  dissolves  in  fused  alkalis  to  form  silicates.  Being  non- 
volatile, it  will  drive  out  most  other  anhydrides  when  heated 


SILICON,  TITANIUM,  BORON  261 

to  a  high  temperature  with  their  salts,  especially  when  the 
silicates  so  formed  are  fusible.  The  following  equations 
illustrate  this  property : 

Na2CO3  +  SiO2  =  Na2SiO3  +  CO2, 
Na2SO4  +  SiO2  =  Na2SiO3  +'SO8. 

Silicic  acids.  Silicon  forms  two  simple  acids,  ortho- 
silicic  acid  (H4SiO4)  and  metasilicic  acid  (H2SiO3).  Ortho- 
silicic  acid  is  formed  as  a  jelly-like  mass  when  orthosilicates 
are  treated  with  strong  acids  such  as  hydrochloric.  On 
attempting  to  dry  this  acid  it  loses  water,  passing  into 
metasilicic  or  common  silicic  acid  : 

H4Si04  -  H2Si03  +  H20. 

Metasilicic  acid  when  heated  breaks  up  into  silica  and 
water,  thus  : 

H2SiO3  =  H2O  +  SiO2.  * 

Salts  of  silicic  acids, — silicates.  A  number  of  salts  of  the  ortho- 
silicic  and  metasilicic  acids  occur  in  nature.  Thus  mica  (KAlSiO4) 
is  a  salt  of  orthosilicic  acid. 

Polysilicic  acids.  Silicon  has  the  power  to  form  a  great 
many  complex  acids  which  may  be  regarded  as  derived 
from  the  union  of  several  molecules  of  the  orthosilicic 
acid,  with  the  loss  of  water.  Thus  we  have 

3  H4Si04  =  H4Si808  +  4  H2O. 

These  acids  cannot  be  prepared  in  the  pure  state,  but 
their  salts  form  many  of  the  crystalline  rocks  in  nature. 
Feldspar,  for  example,  has  the  formula  KAlSi3O8,  and  is  a 
mixed  salt  of  the  acid  H4Si3O8,  whose  formation  is  repre- 
sented in  the  equation  above.  Kaolin  has  the  formula 
Al2Si2O7-2  H2O.  Many  other  examples  will  be  met  in 
the  study  of  the  metals. 


262     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


Glass.  When  sodium  and  calcium  silicates,  together  with  silicon 
dioxide,  are  heated  to  a  very  high  temperature,  the  mixture  slowly 
fuses  to  a  transparent  liquid,  which  on  cooling  passes  into  the  solid 
called  glass.  Instead  of  starting  with  sodium  and  calcium  silicates  it 
is  more  convenient  and  economical  to  heat  sodium  carbonate  (or 

sulphate)  and  lime  with  an  excess  of 
clean  sand,  the  silicates  being  formed 
\B  during  the  heating : 

Na2CO3  +  SiO2  =  Na2SiO3  +  CO2, 
CaO  +  SiO2  =  CaSiO3. 

The  mixture  is  heated  below  the 
fusing  point  for  some  time,  so  that  the 
escaping  carbon  dioxide  may  not  spatter 
the  hot  liquid  ;  the  heat  is  then  increased 
and  the  mixture  kept  in  a  state  of  fusion 
until  all  gases  formed  in  the  reaction 
have  escaped. 

Molding  and  blowing  of  glass.  The 
way  in  which  the  melted  mixture  is 
handled  in  the  glass  factory  depends 
upon  the  character  of  the  article  to  be 
made.  Many  articles,  such  as  bottles,  are  made  by  blowing  the  plastic 
glass  into  hollow  molds  of  the  desired  shape.  The  mold  is  first  opened, 
as  shown  in  Fig.  74.  A  lump  of  plastic  glass  A  on  the  hollow  rod  B 
is  lowered  into  the  mold,  which  is  then  closed  by  the  handles  C.  By 
blowing  into  the  tube  the  glass  is  blown  into  the  shape  of  the  mold. 
The  mold  is  then  opened  and  the  bottle  lifted  out.  The  neck  of  the 
bottle  must  be  cut  off  at  the  proper  place  and  the  sharp  edges 
rounded  off  in  a  flame. 

Other  objects,  such  as  lamp  chimneys,  are  made  by  getting  a  lump 
of  plastic  glass  on  the  end  of  a  hollow  iron  rod  and  blowing  it  into 
the  desired  shape  without  the  help  of  a  mold,  great  skill  being 
required  in  the  manipulation  of  the  glass.  Window  glass  is  made  by 
blowing  large  hollow  cylinders  about  6  ft.  long  and  i|ft.  in  diameter. 
These  are  cut  longitudinally,  and  are  then  placed  in  an  oven  and 
heated  until  they  soften,  when  they  are  flattened  out  into  plates 
(Fig.  75).  Plate  glass  is  cast  into  flat  slabs,  which  are  then  ground 
and  polished  to  perfectly  plane  surfaces. 


FIG.  74 


SILICON,  TITANIUM,  BORON  263 

Varieties  of  glass.  The  ingredients  mentioned  above  make  a  soft, 
easily  fusible  glass.  If  potassium  carbonate  is  substituted  for  the 
sodium  carbonate,  the  glass  is  much  harder  and  less  easily  fused ; 
increasing  the  amount  of  sand  has  somewhat  the  same  effect.  Potas- 
sium glass  is  largely  used  in  making  chemical  glassware,  since  it 
resists  the  action  of  reagents  better  than  the  softer  sodium  glass.  If 
lead  oxide  is  substituted  for  the  whole  or  a  part  of  the  lime,  the  glass 
is  very  soft,  but  has  a  high  index  of  refraction  and  is  valuable  for 
making  optical  instruments  and  artificial  jewels. 

Coloring  of  glass.    Various  substances  fused  along  with  the  glass 
mixture  give  characteristic  colors.      The  amber  color  of  common 
bottles  is  due  to  iron  compounds  in  the  glass;  in  other  cases  iron 
colors  the  glass  green. 
Cobalt  compounds 
color  it  deep  blue; 
those  of  manganese 
give    it    an    amethyst 
tint  and  uranium  com- 
pounds impart  a  pecu-  FlG>  75 
liar   yellowish  green 

color.  Since  iron  is  nearly  always  present  in  the  ingredients,  glass 
is  usually  slightly  yellow.  This  color  can  be  removed  bidding  thp 
proper  amount  of  manganese  dioxide,  for  the  amethyst  rnlnr  of  man- 
ganese and  the"  yellow  of  iron  together  produce  white  li^ht. 

Nature  of  glass.  Glass  is  not  a  definite  chemical  compound  and 
its  composition  varies  between  wide  limits.  Fused  glass  is  really  a 
solution  of  various  silicates,  such  as  those  of  calcium  and  lead,  in 
fused  sodium  or  potassium  silicate.  A  certain  amount  of  silicon 
dioxide  is  also  present.  This  solution  is  then  allowed  to  solidify 
under  such  conditions  of  cooling  that  the  dissolved  substances  do 
not  separate  from  the  solvent.  The  compounds  which  are  used  to 
color  the  glass  are  sometimes  converted  into  silicates,  which  then 
dissolve  in  the  glass,  giving  it  a' uniform  color.  In  other  cases,  as 
in  the  milky  glasses  which  resemble  porcelain  in  appearance,  the 
color  or  opaqueness  is  due  to  the  finely  divided  color  material  evenly 
distributed  throughout  the  glass,  but  not. dissolved  in  it.  Milky  glass 
is  made  by  mixing  calcium  fluoride,  tin  oxide,  or  some  other  insoluble 
substance  in  the  melted  glass.  Copper  or  gold  in  metallic  form 
scattered  through  glass  gives  it  shades  of  red. 


264     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

TITANIUM 

Titanium  is  a  very  widely  distributed  element  in  nature,  being 
found  in  almost  all  soils,  in  many  rocks,  and  even  in  plant  and  animal 
tissues.  It  is  not  very  abundant  in  any  one  locality,  and  it  possesses 
little  commercial  value  save  in  connection  with  the  iron  industry. 
Its  most  common  ore  is  rutile  (TiO2),  which  resembles  silica  in  many 
respects. 

In  both  physical  and  chemical  properties  titanium  resembles 
silicon,  though  it  is  somewhat  more  metallic  in  character.  This 
resemblance  is  most  marked  in  the  acids  of  titanium.  It  not  only 
forms  metatitanic  and  orthotitanic  acids  but  a  great  variety  of  poly- 
titanic  acids  as  well. 

BORON 

Occurrence.  Boron  is  never  found  free  in  nature.  It 
occurs  as  boric  acid  (H3BO3),  and  in  salts  of  polyboric 
acids,  which  usually  have  very  complicated  formulas. 

Preparation  and  properties.  Boron  can  be  prepared  from 
its  oxide  by  reduction  with  magnesium,  exactly  as  in  the 
case  of  silicon.  It  resembles  silicon  very  strikingly  in  its 
properties.  It  occurs  in  several  allotropic  forms,  is  very 
hard  when  crystallized,  and  is  rather  inactive  toward 
reagents.  It  forms  a  hydride,  BH3,  and  combines  directly 
with  the  elements  of  the  chlorine  family.  Boron  fluoride 
(BF3)  is  very  similar  to  silicon  fluoride  in  its  mode  of 
formation  and  chemical  properties. 

Boric  oxide  (B2O3).  Boron  forms  one  well-known  oxide, 
B2O3,  called  boric  anhydride.  It  is  formed  as  a  glassy 
mass  by  heating  boric  acid  to  a  high  temperature.  It 
absorbs  water  very  readily,  uniting  with  it  to  form  boric 
acid  again  :  ^  +  3  HIQ  =  2  H3BO3. 

In  this  respect  it  differs  from  silicon  dioxide,  which  will 
not  combine  directly  with  water. 


SILICON,  TITANIUM,  BORON  265 

Boric  acid  (H3BO3).  This  is  found  in  nature  in  consid- 
erable quantities  and  forms  one  of  the  chief  sources  of 
boron  compounds.  It  is  found  dissolved  in  the  water,  of 
hot  springs  in  some  localities,  particularly  in  Italy.  Being 
volatile  with  steam,  the  vapor  which  escapes  from  these 
springs  has  some  boric  acid  in  it.  It  is  easily  obtained 
from  these  sources  by  condensation  and  evaporation,  the 
necessary  heat  being  supplied  by  other  hot  springs. 

Boric  acid  crystallizes  in  pearly  flakes,  which  are  greasy 
to  the  touch.  In  the  laboratory  it  is  easily  prepared  by 
treating  a  strong,  hot  solution  of  borax  with  sulphuric  acid. 
Boric  acid  being  sparingly  soluble  in  water  crystallizes 
out  on  cooling : 

Na2B4O7  +  5  H2O  +  H2SO4  =  Na2SO4  +  4  H3BO3. 

The  substance  is  a  mild  antiseptic,  and  on  this  account  is 
often  used  in  medicine  and  as  a  preservative  for  canned 
foods  and  milk. 

Metaboric  and  polyboric  acids.  When  boric  acid  is  gently 
heated  it  is  converted  into  metaboric  acid  (HBO2) : 

H3BO3  =  HBO2  +  H2O. 

On  heating  metaboric  acid  to  a  somewhat  higher  tempera- 
ture tetraboric  acid  (H2B4O7)  is  formed  : 

4  HBO2  =  H2B4O7  +  H^p. 

Many  other  complex  acids  of  boron  are  known. 

Borax.  Borax  is  the  sodium  salt  of  tetraboric  acid,  hav- 
ing the  formula  Na2B4O7*io  H2O.  It  is  found  in  some 
arid  countries,  as  southern  California  and  Tibet,  but  is 
now  made  commercially  from  the  mineral  colemanite,  which 
is  the  calcium  salt  of  a  complex  boric  acid.  When  this 
is  treated  with  a  solution  of  sodium  carbonate,  calcium 


266     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

carbonate  is  precipitated  and  borax  crystallizes  from  the 
solution. 

When  heated  borax  at  first  swells  up  greatly,  owing  to 
the  expulsion  of  the  water  of  crystallization,  and  then  melts 
to  a  clear  glass.  This  glass  has  the  property  of  easily  dis- 
solving many  metallic  oxides,  and  on  this  account  borax  is 
used  as  a  flux  in  soldering,  for  the  purpose  of  removing 
from  the  metallic  surfaces  to  be  soldered  the  film  of  oxide 
with  which  they  are  likely  to  be  covered.  These  oxides 
often  give  a  characteristic  color  to  the  clear  borax  glass, 
and  borax  beads  are  therefore  often  used  in  testing  for  the 
presence  of  metals,  instead  of  the  metaphosphoric  acid 
bead  already  described. 

The  reason  that  metallic  oxides  dissolve  in  borax  is  that  borax 
contains  an  excess  of  acid  anhydride,  as  can  be  more  easily  seen  if 
its  formula  is  written  2NaBO2  +  B2O3.  The  metallic  oxide  com- 
bines with  this  excess  of  acid  anhydride,  forming  a  mixed  salt  of 
metaboric  acid. 

Borax  is  extensively  used  as  a  constituent  of  enamels 
and  glazes  for  both  metal  ware  and  pottery.  It  is  also 
used  as  a  flux  in  soldering  and  brazing,  and  in  domestic 
ways  it  serves  as  a  mild  alkali,  as  a  preservative  for  meats, 
and  in  a  great  variety  of  less  important  applications. 

EXERCISES 

1.  Account  for  the  fact  that  a  solution  of  borax  in  water  is  alkaline. 

2.  What  weight  of  water  of  crystallization  does  i   kg.  of  borax 
contain  ? 

3.  When  a  concentrated  solution  of  borax  acts  on  silver  nitrate  a 
borate  of  silver  is  formed.    If  the  solution  of  borax  is  dilute,  however, 
an  hydroxide  of  silver  forms.    Account  for  this  difference  in  behavior. 


CHAPTER   XXII 
THE  METALS 

The  metals.  The  elements  which  remain  to  be  consid- 
ered are  known  collectively  as  the  metals.  They  are  also 
called  the  base-forming  elements,  since  their  hydroxides  are 
bases.  *V  metal  may  therefore  be  defined  as  an  element 
whose  hydroxide  is  a  base.\  When  a  base  dissolves  in 
water  the  hydroxyl  groups  form  the  anions,  while  the 
metallic  element  forms  the  cations.  From  this  standpoint 
a  metal  can  be  defined  as  an  element  capable  of  forming 
simple  cations  in  solution. 

The  distinction  between  a  metal  and  a  non-metal  is  not 
a  very  sharp  one,  since  the  hydroxides  of  a  number*  of  ele- 
ments act  as  bases  under  some  conditions  and  as  acids 
under  others.  We  have  seen  that  antimony  is  an  element 
of  this  kind. 

Occurrence  of  metals  in  nature.  A  few  of  the  metals  are 
found  in  nature  in  the  free  state.  Among  these  are  gold, 
platinum,  and  frequently  copper.  They  are  usually  found 
combined  with  other  elements  in  the 'form  of  oxides  or 
salts  of  various  acids.  Silicates,  carbonates,  sulphides,  and 
sulphates  are  the  most  abundant  salts.  All  inorganic  sub- 
stances occurring  in  nature,  whether  they  contain  a  metal 
or  not,  are  called  minerals.  Those  minerals  from  which  a 
useful  substance  can  be  extracted  are  called  ores  of  the 
substance.  These  two  terms  are  most  frequently  used  in 
connection  with  the  metals. 

267 


268     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Extraction  of  metals,  —  metallurgy.  The  process  of  ex- 
tracting a  metal  from  its  ores  is  called  the  metallurgy  of 
the  metal.  The  metallurgy  of  each  metal  presents  pecul- 
iarities of  its  own,  but  there  are  several  methods  of  general 
application  which  are  very  frequently  employed. 

i.  Reduction  of  an  oxide  with  carbon.  Many  of  the 
metals  occur  in  nature  in  the  form  of  oxides.  When  these 
oxides  are  heated  to  a  high  temperature  with  carbon  the 
oxygen  combines  with  it  and  the  metal  is  set  free.  Iron, 
for  example,  occurs  largely  in  the  form  of  the  oxide  Fe2O3. 
When  this  is  heated  with  carbon  the  reaction  expressed  in 
the  following  equation  takes  place  : 

Fe2O3  -f  3  C  =  2  Fe  +  3  CO. 

Many  ores  other  than  oxides  may  be  changed  into  oxides 
which  can  then  be  reduced  by  carbon.  The  conversion  of 
such  ores  into  oxides  is  generally  accomplished  by  heating, 
and  this  process  is  called  roasting.  Many  carbonates  and 
hydroxides  decompose  directly  into  the  oxide  on  heating. 
Sulphides,  on  the  other  hand,  must  be  heated  in  a  current 
of  air,  the  oxygen  of  the  air  entering  into  the  reaction. 
The  following  equations  will  serve  to  illustrate  these  changes 
in  the  case  of  the  ores  of  iron  : 

FeCO3  =  FeO  4-  CO2, 


2  FeS2  +  1  1  O  =  Fe2O3  +  4  SO2. 

2.  Reduction  of  an  oxide  with  aluminium.  Not  all  oxides, 
however,  can  be  reduced  by  carbon.  In  such  cases  alu- 
minium may  be  used.  Thus  chromium  may  be  obtained  in 
accordance  with  the  following  equation  : 

Cr2O3  +  2  Al  =  2  Cr  +  A12O3. 


THE  METALS  269 

This  method  is  a  comparatively  new  one,  having  been 
brought  into  use  by  the  German  chemist  Goldschmidt ; 
hence  it  is  sometimes  called  the  Goldschmidt  method. 

3.  Electrolysis.  In  recent  years  increasing  use  is  being 
made  of  the  electric  current  in  the  preparation  of  metals. 
In  some  cases  the  separation  of  the  metal  from  its  com- 
pounds is  accomplished  by  passing  the  current  through  a 
solution  of  a  suitable  salt  of  the  metal,  the  metal  usually 
being  deposited  upon  the  cathode.  In  other  cases  the 
current  is  passed  through  a  fused  salt  of  the  metal,  the 
chloride  being  best  adapted  to  this  purpose. 

Electro-chemical  industries.  Most  of  the  electro-chemical 
industries  of  the  country  are  carried  on  where  water  power 
is  abundant,  since  this  furnishes  the  cheapest  means  for 
the  generation  of  electrical  energy.  Niagara  Falls  is  the 
most  important  locality  in  this  country  for  such  industries, 
and  many  different  electro-chemical  products  are  manu- 
factured there.  Some  industries  depend  upon  electrolytic 
processes,  while  in  others  the  electrical  energy  is  used 
merely  as  a  source  of  heat  in  electric  furnaces. 

Preparation  of  compounds  of  the  metals.  Since  the  com- 
pounds of  the  metals  are  so  numerous  and  varied  in  char- 
acter, there  are  many  ways  of  preparing  them.  In  many 
cases  the  properties  of  the  substance  to  be  prepared,  or  the 
material  available  for  its  preparation*  suggest  a  rather 
unusual  way.  There  are,  however,  a  number  of  general 
principles  which  are  constantly  applied  in  the  preparation 
of  the  compounds  of  the  metals,  and  a  clear  understanding 
of  them  will  save  much  time  and  effort  in  remembering  the 
details  in  any  given  case.  The  most  important  of  these 
general  methods  for  the  preparation  of  compounds  are 
the  following : 


270     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

1.  By    direct   union    of  two   elements.    This   is   usually 
accomplished  by  heating  the  two  elements  together.    Thus 
the  sulphides,  chlorides,  and  oxides  of  a  metal  can  generally 
be  obtained  in  this  way.    The  following  equations  serve  as 
examples  of  this  method  : 

Fe  +  S  =  FeS, 

Mg  +  O  =  MgO, 

Cu  +  2  Cl  =  CuCl2. 

2.  By  the  decomposition  of  a  compound.    This   decom- 
position may  be  brought  about  either  by  heat  alone  or  by 
the  combined  action  of  heat  and  a  reducing  agent.    Thus 
when   the  nitrate  of  a  metal   is  heated  the  oxide  of  the 
metal  is  usually   obtained.    Copper   nitrate,  for  example, 
decomposes  as  follows  : 

Cu(NO3)2  =  CuO  +  2  NO2  +  O. 
Similarly  the  carbonates  of  the  metals  yield  oxides,  thus  : 

CaCO3  =  CaO  +  CO2. 

Most  of  the  hydroxides  form  an  oxide  and  water  when 
heated  :  2  A1(OH)3  =  A12O3  +  3  H2O. 

When  heated  with  carbon,  sulphates  are  reduced  to  sul- 
phides,  thus  :   BaS0  +  2  c  =  BaS  + 


3  .  Methods  based  on  equilibrium  in  solution.  In  the  prepa- 
ration of  compounds  the  first  requisite  is  that  the  reactions 
chosen  shall  be  of  such  a  kind  as  will  go  on  to  completion. 
In  the  chapter  on  chemical  equilibrium  it  was  shown  that 
reactions  in  solution  may  become  complete  in  either  of 
three  ways  :  (i)  a  gas  may  be  formed  which  escapes  from 
solution  ;  (2)  an  insoluble  solid  may  be  formed  which  pre- 
cipitates ;  (3)  two  different  ions  may  combine  to  form 


THE  METALS  271 

undissociated  molecules.  By  the  judicious  selection  of 
materials  these  principles  may  be  applied  to  the  preparation 
of  a  great  variety  of  compounds,  and  illustrations  of  such 
methods  will  very  frequently  be  found  in  the  subsequent 
pages. 

4.  By  fusion  methods.    It  sometimes  happens  that  sub- 
stances  which  are   insoluble   in  water  and  in  acids,  and 
which  cannot  therefore  be  brought  into  double  decomposi- 
tion in  the  usual  way,  are  soluble  in  other  liquids,  and  when 
dissolved  in  them  can  be  decomposed  and  converted  into 
other  desired  compounds.     Thus  barium  sulphate   is  not 
soluble  in  water,  and  sulphuric  acid,  being  less  volatile  than 
most  other  acids,  cannot  easily  be  driven  out  from  this  salt. 
When  brought  into  contact  with  melted  sodium  carbonate, 
however,  it  dissolves  in  it,  and  since  barium  carbonate  is 
insoluble  in   melted   sodium   carbonate,  double   decompo- 
sition takes  place  : 

Na2CO3  +  BaSO4  =  BaCO3  +  Na2SO4. 

On  dissolving  the  cooled  mixture  in  water  the  sodium  sul- 
phate formed  in  the  reaction,  together  with  any  excess  of 
sodium  carbonate  which  may  be  present,  dissolves.  The 
barium  carbonate  can  then  be  filtered  off  and  converted 
into  any  desired  salt  by  the  processes  already  described. 

5 .  By  .the  action  of  metals  on  salts  of  other  metals.    When 
a  strip  of  zinc  is  placed  in  a  solution  of  a  copper  salt  the 
copper  is  precipitated  and  an  equivalent  quantity  of  zinc 
passes  into  solution  : 

Zn  +  CuSO4  -  Cu  +  ZnSO4.  ' 

In  like  manner  copper  will  precipitate  silver  from  its  salts : 
Cu  +  Ag2SO4  =  2  Ag  +  CuSO4. 


272     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

It  is  possible  to  tabulate  the  metals  in  such  a  way  that  any 
one  of  them  in  the  table  will  precipitate  any  one  following 
it  from  its  salts.  The  following  is  a  list  of  some  of  the 
commoner  metals  arranged  in  this  way : 

Zinc  Lead  Mercury 

Iron  Copper  Silver 

Tin  Bismuth.  Gold 

According  to  this  table  copper  will  precipitate  bismuth, 
mercury,  silver,  or  gold  from  their  salts,  and  will  in  turn 
be  precipitated  by  zinc,  iron,  tin,  or  lead.  Advantage  is 
taken  of  this  principle  in  the  purification  of  some  of  the 
metals,  and  occasionally  in  the  preparation  of  metals  and 
their  compounds. 

Important  insoluble  compounds.  Since  precipitates  play 
so  important  a  part  in  the  reactions  which  substances 
undergo,  as  well  as  in  the  preparation  of  many  chemical 
compounds,  it  is  important  to  know  what  substances  are 
insoluble.  Knowing  this,  we 'can  in  many  cases  predict 
reactions  under  certain  conditions,  and  are  assisted  in 
devising  ways  to  prepare  desired  compounds.  While  there 
is  no  general  rule  which  will  enable  one  to  foretell  the  solu- 
bility of  any  given  compound,  nevertheless  a  few  general 
statements  can  be  made  which  will  be  of  much  assistance. 

1.  Hydroxides.    All  hydroxides  are  insoluble  save  those 
of  ammonium,   sodium,   potassium,  calcium,  barium,  and 
strontium. 

2.  Nitrates.    All  nitrates  are  soluble  in  water. 

3.  Chlorides.    All  chlorides  are  soluble  save  silver  and 
mercurous  chlorides.   (Lead  chloride  is  but  slightly  soluble.) 

4.  Sulphates.    All   sulphates  are  soluble  save  those  of 
barium,  strontium,  and  lead.    (Sulphates  of  silver  and  cal- 
cium are  only  moderately  soluble.) 


THE  METALS  273 

5.  Sulphides.    All  sulphides  are  insoluble  save  those  of 
ammonium,  sodium,  and  potassium.    The  sulphides  of  cal- 
cium, barium,  strontium,  and  magnesium  are  insoluble  in 
water,  but  are  changed  by  hydrolysis  into  acid  sulphides 
which  are  soluble.     On  this  account  they  cannot  be  pre- 
pared by  precipitation. 

6.  Carbonates,  phosphates,  and  silicates.    All  normal  car- 
bonates, phosphates,  and  silicates  are  insoluble  save  those 
of  ammonium,  sodium   and  potassium. 

EXERCISES 

1.  Write  equations  representing  four  different  ways  for  preparing 
Cu(N03)2. 

2.  Write  equations  representing  six  different  ways  for  preparing 
ZnSO4. 

3.  Write  equations  for  two  reactions  to  illustrate  each  of   the 
three  ways  in  which  reactions  in  solutions  may  become  complete. 

4.  Give  one  or  more  methods  for  preparing  each  of  the  following 
compounds:   CaCl2,  PbCl2,  BaSO4,  CaCO3,  (NH4)2S,  Ag2S,  PbO, 
Cu(OH)2  (for  solubilities,  see  last  paragraph  of  chapter).     State  in 
each  case  the  general  principle  involved  in  the  method  of  preparation 
chosen. 


CHAPTER   XXIII 
THE  ALKALI  METALS 


SYMBOL 

ATOMIC 
WEIGHT 

DENSITY 

MELTING 
POINT 

FIRST  PREPARED 

Lithium  .     .     . 

Li 

7-03 

o-59 

186.° 

Davy        1820 

Sodium   .     .     . 

Na 

23-05 

0.97 

97-6° 

"            1807 

Potassium    .     . 

K 

39-15 

0.87 

62.5° 

1807 

Rubidium    .     . 

Rb 

85.5 

1.52 

38-5° 

Bunsen     1861 

Caesium  .     .     . 

Cs 

132.9 

1.88 

26.5° 

1860 

The  family.  The  metals  listed  in  the  above  table  con- 
stitute the  even  family  in  Group  I  in  the  periodic  arrange- 
ment of  the  elements,  and  therefore  form  a  natural  family. 
The  name  alkali  metals  is  commonly  applied  to  the 
family  for  the  reason  that  the  hydroxides  of  the  most 
familiar  members  of  the  family,  namely  sodium  and 
potassium,  have  long  been  called  alkalis. 

1.  Occurreme.    While  none  of  these  metals  occur  free 
in  nature,  their  compounds  are   very  widely  distributed, 
being  especially  abundant  in  sea  and  mineral  waters,  in 
salt  beds,  and  in  many  rocks.    Only  sodium  and  potassium 
occur  in  abundance,  the  others  being  rarely  found  in  any 
considerable .  quantity. 

2.  Preparation.    The  metals  are  most  conveniently  pre- 
pared  by  the  electrolysis    of    their  fused    hydroxides    or 
chlorides,  though  it  is  possible  to  prepare  them  by  reducing 
their  oxides  or  carbonates  with  carbon. 

274 


THE  ALKALI  METALS  275 

3.  Properties.    They  are  soft,  light  metals,  having  low 
melting  points  and  small  densities,  as  is  indicated  in  the 
table.    Their    melting   points    vary    inversely    with    their 
atomic  weights,  while  their  densities  (sodium    excepted) 
vary  directly  with  these.    The  pure  metals  have  a  silvery 
luster  but  tarnish  at  once  when  exposed  to  the  air,  owing 
to  the  formation  of  a  film  of  oxide  upon  the  surface  of  the 
metal.    They  are  therefore  preserved  in  some  liquid,  such 
as  coal  oil,  which  contains  no  oxygen.     Because  of  their 
strong   affinity   for   oxygen    they    decompose  water  with 
great  ease,  forming  hydroxides   and  liberating  hydrogen 
in  accordance  with  the  equation 

M  +  H2O  =  MOH  +  H, 

where  M  stands  for  any  one  of  these  metals.  These 
hydroxides  are  white  solids ;  they  are  readily  soluble  in 
water  and  possess  very  strong  basic  properties.  These 
bases  are  nearly  equal  in  strength,  that  is,  they  all  dis- 
sociate in  water  to  about  the  same  extent. 

4.  Compounds.    The  alkali  metals  almost  always  act  as 
univalent  elements  in  the  formation  of   compounds,  the 
composition  of  which  can  be  represented  by  such  formulas 
as  MH,  MCI,  MNO3,  M2SO4,  M3PO4.    These  compounds, 
when  dissolved  in  water,  dissociate  in  such  a  way  as  to 
form  simple,  univalent  metallic  ions  which  are  colorless. 
With  the  exception  of  lithium  these  metals  form  very  few 
insoluble  compounds,  so  that  it  is  not  often  that  precipi- 
tates containing    them    are   obtained.     Only  sodium  and 
potassium  will  be  studied  in  detail,  since  the  other  metals 
of  the  family  are  of  relatively  small  importance. 

The  compounds  of  sodium  and  potassium  are  so  similar 
in  properties  that  they  can  be   used  interchangeably  for 


276     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

most  purposes.  Other  things  being  equal,  the  sodium 
compounds  are  prepared  in  preference  to  those  of  potas- 
sium, since  they  are  cheaper.  When  a  given  sodium 
compound  is  deliquescent,  or  is  so  soluble  that  it  is  diffi- 
cult to  purify,  the  corresponding  potassium  compound  is 
prepared  in  its  stead,  provided  its  properties  are  more 
desirable  in  these  respects. 

SODIUM 

Occurrence  in  nature.  Large  deposits  of  sodium  chloride 
have  been  found  in  various  parts  of  the  world,  and  the 
water  of  the  ocean  and  of  many  lakes  and  springs  contains 
notable  quantities  of  it.  The  element  also  occurs  as  a  con- 
stituent of  many  rocks  and  is  therefore  present  in  the 
soil  formed  by  their  disintegration.  The  mineral  cryolite 
(Na3AlF6)  is  an  important  substance,  and  the  nitrate,  car- 
bonate, and  borate  also  occur  in  nature. 

Preparation.  In  1807  Sir  Humphry  Davy  succeeded 
in  preparing  very  small  quantities  of  metallic  sodium  by 
the  electrolysis  of  the  fused  hydroxide.  On  account  of 
the  cost  of  electrical  energy  it  was  for  many  years  found 
more  economical  to  prepare  it  by  reducing  the  carbonate 
with  carbon  in  accordance  with  the  following  equation : 

Na2CO3  +  2C  =  2Na+3  CO. 

The  cost  of  generating  the  electric  current  has  been  dimin- 
ished to  such  an  extent,  however,  that  it  is  now  more  eco- 
nomical to  prepare  sodium  by  Davy's  original  method, 
namely,  by  the  electrolysis  of  the  fused  hydroxide  or  chlo- 
ride. When  the  chloride  is  used  the  process  is  difficult  to 
manage,  owing  to  the  higher  temperature  required  to  keep 
the  electrolyte  fused,  and  because  of  the  corroding  action  of 
the  fused  chloride  upon  the  containing  vessel. 


THE  ALKALI  METALS 


277 


Technical  preparation.  The  sodium  hydroxide  is  melted  in  a  cylin- 
drical iron  vessel  (Fig.  76)  through  the  bottom  of  which  rises  the 
cathode  K.  The  anodes  A,  several  in  number,  are  suspended  around 
the  cathode  from  above.  A  cylindrical  vessel  C  floats  in  the  fused 
alkali  directly  over  the  cathode,  and  under  this  cap  the  sodium  and 
hydrogen  liberated  at  the  cathode  collect.  The  hydrogen  escapes  by 
lifting  the  cover,  and  the  sodium,  pro- 
tected from  the  air  by  the  hydrogen,  is 
skimmed  or  drained  off  from  time  to  + 
time.  Oxygen  is  set  free  upon  the  an- 
ode and  escapes  into  the  air  through 
the  openings  O  without  coming  into 
contact  with  the  sodium  of  hydrogen. 
This  process  is  carried  on  extensively 
at  Niagara  Falls. 

Properties.  Sodium  is  a  silver- 
white  metal  about  as  heavy  as 
water,  and  so  soft  that  it  can 
be  molded  easily  by  the  fingers 
or  pressed  into  wire.  It  is  very  FlG  76 

active  chemically,  combining 

with  most  of  the  non-metallic  elements,  such  as  oxygen  and 
chlorine,  with  great  energy.  It  will  often  withdraw  these 
elements  from  combination  with  other  elements,  and  is  thus 
able  to  decompose  water  and  the  oxides  and  chlorides  of 
many  metals. 

Sodium  peroxide  (NaO).  Since  sodium  is  a  univalent 
element  we  should  expect  it  to  form  an  oxide  of  the 
formula  Na2O.  While  such  an  oxide  can  be  prepared,  the 
peroxide  (NaO)  is  much  better  known.  It  is  a  yellowish- 
white  powder  made  by  burning  sodium  in  air.  Its  chief 
use  is  as  an  oxidizing  agent.  When  heated  with  oxidizable 
substances  it  gives  up  a  part  of  its  oxygen,  as  shown  in 

the  equation  AT  /^      AT    r\  ,  r\ 

2  NaO  =  Na2O  4-  O. 


278     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Water  decomposes  it  in  accordance  with  the  equation 
2  NaO  +  2  H2O  =  2  NaOH  +  H2O2. 

Acids  act  readily  upon  it,  forming  a  sodium  salt  and 
hydrogen  peroxide  : 

2  NaO  +  2  HC1  =  2  NaCl  +  H2O2. 

In  these  last  two  reactions  the  hydrogen  dioxide  formed 
may  decompose  into  water  and  oxygen  if  the  temperature 
is  allowed  to  rise:  H2O2  =  H2O  +  O. 

Peroxides.  It  will  be  remembered  that  barium  dioxide  (BaO2) 
yields  hydrogen  dioxide  when  treated  with  acids,  and  that  manganese 
dioxide  gives  up  oxygen  when  heated  with  sulphuric  acid.  Oxides 
which  yield  either  hydrogen  dioxide  or  oxygen  when  treated  with 
water  or  an  acid  are  called  peroxides. 

Sodium  hydroxide  (caustic  soda)  (NaOH).  i.  Prepa- 
ration. Sodium  hydroxide  is  prepared  commercially  by 
several  processes. 

(a)  In  the  older  process,  still  in  extensive  use,  sodium 
carbonate  is  treated  with  calcium  hydroxide  suspended  in 
water.  Calcium  carbonate  is  precipitated  according  to  the 
equation 

Na2CO3  +  Ca(OH)2  -  CaCO3  +  2  NaOH. 

The  dilute  solution  of  sodium  hydroxide,  filtered  from  the 
calcium  carbonate,  is  evaporated  to  a  paste  and  is  then 
poured  into  molds  to  solidify.  It  is  sold  in  the  form  of 
slender  sticks. 

(b}  The  newer  methods  depend  upon  the  electrolysis  of 
sodium  chloride.  In  the  Castner  process  a  solution  of  salt 
is  electrolyzed,  the  reaction  being  expressed  as  follows: 

NaCl  +  H90  =  NaOH  +  H  +  CL 


THE  ALKALI  METALS 


279 


The  chlorine  escapes  as  a  gas,  and  by  an  ingenious 
mechanical  device  the  sodium  hydroxide  is  prevented 
from  mixing  with  the  salt  in  the  solution. 

In  the  Acker  process  the  electrolyte  is  fused  sodium 
chloride.  The  chlorine  is  evolved  as  a  gas  at  the  anode, 
while  the  sodium  alloys  with  the  melted  lead  which  forms 
the  cathode.  When  this  alloy  is  treated  with  water  the 
following  reaction  takes  place  : 

Na  +  H2O  =  NaOH  +  H. 

Technical  process.  A  sketch  of  an  Acker  furnace  is  represented  in 
Fig.  77.  The  furnace  is  an  irregularly  shaped  cast-iron  box,  divided 


FIG. 77 


into  three  compartments,  A,  B,  and  C.  Compartment  A  is  lined 
with  magnesia  brick.  Compartments  B  and  C  are  filled  with  melted 
lead,  which  also  covers  the  bottom  of  A  to  a  depth  of  about  an  inch. 
Above  this  layer  in  A  is  fused  salt,  into  which  dip  carbon  anodes  D. 
The  metallic  box  and  melted  lead  is  the  cathode. 

When  the  furnace  is  in  operation  chlorine  is  evolved  at  the 
anodes,  and  is  drawn  away  through  a  pipe  (not  represented)  to 
the  bleaching-powder  chambers.  Sodium  is  set  free  at  the  surface  of 
the  melted  lead  in  A,  and  at  once  alloys  with  it.  Through  the  pipe 
E  a  powerful  jet  of  steam  is  driven  through  the  lead  in  B  upwards 


280     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

into  the  narrow  tube  F.    This  forces  the  lead  alloy  up  through  the 
tube  and  over  into  the  chamber  G. 

In  this  process  the  steam  is  decomposed  by  the  sodium  in  the 
alloy,  forming  melted  sodium  hydroxide  and  hydrogen.  The  melted 
lead  and  sodium  hydroxide  separate  into  two  layers  in  G,  and  the 
sodium  hydroxide,  being  on  top,  overflows  into  tanks  from  which  it  is 
drawn  off  and  packed  in  metallic  drums.  The  lead  is  returned  to  the 
other  compartments  of  the  furnace  by  a  pipe  leading  from  H  to  /. 
Compartment  C  serves  merely  as  a  reservoir  for  excess  of  melted  lead. 

2.  Properties.  Sodium  hydroxide  is  a  white,  crystalline, 
brittle  substance  which  rapidly  absorbs  water  and  carbon 
dioxide  from  the  air.  As  the  name  (caustic  soda)  indicates, 
it  is  a  very  corrosive  substance,  having  a  disintegrating 
action  on  most  animal  and  vegetable  tissues.  It  is  a  strong 
base.  It  is  used  in  a  great  many  chemical  industries,  and 
under  the  name  of  lye  is  employed  to  a  small  extent  as  a 
cleansing  agent  for  household  purposes. 

Sodium  chloride  (common  salt]  (NaCl).  i.  Preparation. 
Sodium  chloride,  or  common  salt,  is  very  widely  distributed 
in  nature.  Thick  strata,  evidently  deposited  at  one  time 
by  the  evaporation  of  salt  water,  are  found  in  many  places. 
In  the  United  States  the  most  important  localities  for  salt 
are  New  York,  Michigan,  Ohio,  and  Kansas.  Sometimes 
the  salt  is  mined,  especially  if  it  is  in  the  pure  form  called 
rock  salt.  More  frequently  a  strong  brine  is  pumped  from 
deep  wells  sunk  into  the  salt  deposit,  and  is  then  evaporated 
in  large  pans  until  the  salt  crystallizes  out.  The  crystals 
are  in  the  form  of  small  cubes  and  contain  no  water  of 
crystallization ;  s»me  water  is,  however,  held  in  cavities 
in  the  crystals  and  causes  the  salt  to  decrepitate  when 
heated. 

2.  Uses.  Since  salt  is  so  abundant  in  nature  it  forms 
the  starting  point  in  the  preparation  of  all  compounds 


THE  ALKALI  METALS  281 

containing  either  sodium  or  chlorine.  This  includes  many 
substances  of  the  highest  importance  to  civilization,  such  as 
soap,  glass,  hydrochloric  acid,  soda,  and  bleaching  powder. 
Enormous  quantities  of  salt  are  therefore  produced  each 
year.  Small  quantities  are  essential  to  the  life  of  man  and 
animals.  Pure  salt  does  not  absorb  moisture ;  the  fact 
that  ordinary  salt  becomes  moist  in  air  is  not  due  to  a 
property  of  the  salt,  but  to  impurities  commonly  occurring 
in  it,  especially  calcium  and  magnesium  chlorides. 

Sodium  sulphate  (Glaubers  salt},  (Na2SO4-  ioH2O). 
This  salt  is  prepared  by  the  action  of  sulphuric  acid  upon 
sodium  chloride,  hydrochloric  acid  being  formed  at  the 
same  time  : 

2  NaCl  +  H2SO4  =  Na2SO4  +  2  HC1. 

Some  sodium  sulphate  is  prepared  by  the  reaction  repre- 
sented in  the  equation 

MgSO4  +  2  NaCl  =  Na2SO4  +  MgCl2. 

The  magnesium  sulphate  required  for  this  reaction  is  ob- 
tained in  large  quantities  in  the  manufacture  of  potassium 
chloride,  and  being  of  little  value  for  any  other  purpose  is 
used  in  this  way.  The  reaction  depends  upon  the  fact  that 
sodium  sulphate  is  the  least  soluble  of  any  of  the  four  fac- 
tors in  the  equation,  and  therefore  crystallizes  out  when 
hot,  saturated  solutions  of  magnesium  sulphate  and  sodium 
chloride  are  mixed  together  and  the  resulting  mixture 
cooled. 

Sodium  sulphate  forms  large  efflorescent  crystals.  The 
salt  is  extensively  used  in  the  manufacture  of  sodium  car- 
bonate and  glass.  Small  quantities  are  used  in  medicine. 

Sodium  sulphite  (Na2SO3-  7  H2O).  Sodium  sulphite  is 
prepared  by  the  action  of  sulphur  dioxide  upon  solutions 


282     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

of  sodium  hydroxide,  the  reaction  being  analogous  to  the 
action  of  carbon  dioxide  upon  sodium  hydroxide.  Like  the 
carbonate,  the  sulphite  is  readily  decomposed  by  acids : 

Na2S03  +  2  HC1  =  2  NaCl  +  H2O  +  SO2. 

Because  of  this  reaction  sodium  sulphite  is  used  as  a 
convenient  source  of  sulphur  dioxide.  It  is  also  used  as  a 
disinfectant  and  a  preservative. 

Sodium  thiosulphate  (Jiyposulphite  of  soda  or  "hypo") 
(Na2S2O3.  5  H2O).  This  salt,  commonly  called  sodium 
hyposulphite,  or  merely  hypo,  is  made  by  boiling  a  solu- 
tion of  sodium  sulphite  with  sulphur : 

Na2SO3  +  S  =  Na2S2O3. 

It  is  used  in  photography  and  in  the  bleaching  industry, 
to  absorb  the  excess  of  chlorine  which  is  left  upon  the 
bleached  fabrics. 

Thio  compounds.  The  prefix  "  thio  "  means  sulphur.  It  is  used  to 
designate  substances  which  may  be  regarded  as  derived  from  oxygen 
compounds  by  replacing  the  whole  or  a  part  of  their  oxygen  with 
sulphur.  The  thiosulphates  may  be  regarded  as  sulphates  in  which 
one  atom  of  oxygen  has  been  replaced  by  an  atom  of  sulphur.  This 
may  be  seen  by  comparing  the  formula  Na2SO4  (sodium  sulphate) 
with  the  formula  Na2S2O3  (sodium  thiosulphate). 

Sodium  carbonate  (sal  soda)  (Na2CO3-  ioH2O).  There 
are  two  different  methods  now  employed  in  the  manufac- 
ture of  this  important  substance. 

I.  Le  Blanc  process.  This  older  process  involves  sev- 
eral distinct  reactions,  as  shown  in  the  following  equations. 

(a)  Sodium  chloride  is  first  converted  into  sodium  sul- 
phate : 

2  NaCl  +  H2S04  =  Na2S04  +  2  HC1. 


THE  ALKALI  METALS  283 

(b)  The  sodium  sulphate  is  next  reduced  to  sulphide  by 
heating  it  with  carbon  : 

Na2SO4  +  2  C  =  Na2S  +  2  CO2. 

(c)  The  sodium   sulphide  is  then  heated  with  calcium 
carbonate,  when  double  decomposition  takes  place: 

Na2S  +  CaCO3  =  CaS  +  Na2CO3. 

Technical  preparation  of  sodium  carbonate.  In  a  manufacturing 
plant  the  last  two  reactions  take  place  in  one  process.  Sodium  sul- 
phate, coal,  and  powdered  limestone  are  heated  together  to  a  rather 
high  temperature.  The  coal  reduces  the  sulphate  to  sulphide,  which 
in  turn  reacts  upon  the  calcium  carbonate.  Some  limestone  is  decom- 
posed by  the  heat,  forming  calcium  oxide.  When  treated  with  water 
the  calcium  oxide  is  changed  into  hydroxide,  and  this  prevents  the 
water  from  decomposing  the  insoluble  calcium  sulphide. 

The  crude  product  of  the  process  is  a  hard  black  cake  called 
black  ash.  On  digesting  this  mass  with  water  the  sodium  carbonate 
passes  into  solution.  The  pure  carbonate  is  obtained  by  evaporation 
of  this  solution,  crystallizing  from  it  in  crystals  of  the  formula 
Na2CO3  •  10  H2O.  Since  over  60%  of  this  salt  is  water,  the  crystals  are 
sometimes  heated  until  it  is  driven  off.  The  product  is  called  calcined 
soda,  and  is,  of  course,  more  valuable  than  the  crystallized  salt. 

2.  Solvay  process.  This  more  modern  process  depends 
upon  the  reactions  represented  in  the  equations 

NaCl  +  NH4HCO3  =  NaHCO3  +  NH4C1, 

2  NaHCO3  =  Na2CO3  +  H2O  +  CO2. 

The  reason  the  first  reaction  takes  place  is'  that  sodium 
hydrogen  carbonate  is  sparingly  soluble  in  water,  while  the 
other  compounds  are  freely  soluble.  When  strong  solutions 
of  sodium  chloride  and  of  ammonium  hydrogen  carbonate 
are  brought  together  the  sparingly  soluble  sodium  hydrogen 
carbonate  is  precipitated.  This  is  converted  into  the  normal 
carbonate,  by  heating,  the  reaction  being  represented  in  the 
second  equation. 


284     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Technical  preparation.  In  the  Solvay  process  a  very  concentrated 
solution  of  salt  is  first  saturated  with  ammonia  gas,  and  a  current  of 
carbon  dioxide  is  then  conducted  into  the  solution.  In  this  way 
ammonium  hydrogen  carbonate  is  formed : 

NH3  +  H20  +  C02  =  NH4HC03. 

This  enters  into  double  decomposition  with  the  salt,  as  shown  in  the 
first  equation  under  the  Solvay  process.  After  the  sodium  hydrogen 
carbonate  has  been  precipitated  the  mother  liquors  containing  ammo- 
nium chloride  are  treated  with  lime : 

2  NH4C1  +  CaO  =  CaCl2  +  2  NH3  +  H2O. 
The  lime  is  obtained  by  burning  limestone : 

CaCO3  =  CaO  +  CO2. 

The  ammonia  and  carbon  dioxide  evolved  in  the  latter  two  reactions 
are  used  in  the  preparation  of  an  additional  quantity  of  ammonium 
hydrogen  carbonate.  It  will  thus  be  seen  that  there  is  no  loss  of 
ammonia.  The  only  materials  permanently  used  up  are  calcium  car- 
bonate and  salt,  while  the  only  waste  product  is  calcium  chloride. 

Historical.  In  former  times  sodium  carbonate  was  made  by  burn- 
ing seaweeds  and  extracting  the  carbonate  from  their  ash.  On  this 
account  the  salt  was  called  soda  ash,  and  the  name  is  still  in  common 
use.  During  the  French  Revolution  this  supply  was  cut  off,  and  in 
behalf  of  the  French  government  Le  Blanc  made  a  study  of  methods 
of  preparing  the  carbonate  directly  from  salt.  As  a  result  he  devised 
the  method  which  bears  his  name,  and  which  was  used  exclusively 
for  many  years.  It  has  been  replaced  to  a  large  extent  by  the  Solvay 
process,  which  has  the  advantage  that  the  materials  used  are  inex- 
pensive, and  that  the  ammonium  hydrogen  carbonate  used  can  be 
regenerated  from  the  products  formed  in  the  process.  Much  expense 
is  also  saved  in  fuel,  and  the  sodium  hydrogen  carbonate,  which  is 
the  first  product  of  the  process,  has  itself  many  commercial  uses. 
The  Le  Blanc  process  is  still  used,  however,  since  the  hydrochloric 
acid  generated  is  of  value. 

By-products.  The  substances  obtained  in  a  given  process,  aside 
from  the  main  product,  are  called  the  by-products.  The"  success  of 
many  processes  depends  upon  the  value  of  the  by-products  formed. 


THE  ALKALI  METALS  285 

Thus  hydrochloric  acid,  a  by-product  in  the  Le  Blanc  process,  is 
valuable  enough  to  make  the  process  pay,  even  though  sodium  car- 
bonate can  be  made  cheaper  in  other  ways. 

Properties  of  sodium  carbonate.  Sodium  carbonate  forms 
large  crystals  of  the  formula  Na2CO3  •  10  H2O.  It  has  a 
mild  alkaline  reaction  and  is  used  for  laundry  purposes 
under  the  name  of  washing  soda.  Mere  mention  of  the 
fact  that  it  is  used  in  the  manufacture  of  glass,  soap,  and 
many  chemical  reagents  will  indicate  its  importance  in 
the  industries.  It  is  one  of  the  few  soluble  carbonates. 

Sodium  hydrogen  carbonate  (bicarbonate  of  soda) 
(NaHCO3).  This  salt,  commonly  called  bicarbonate  of  soda, 
or  baking  soda,  is  made  by  the  Solvay  process,  as  explained 
above,  or  by  passing  carbon  dioxide  into  strong  solutions 
of  sodium  carbonate  : 

Na2CO3  +  H2O  +  CO2  =  2  NaHCO3. 

The  bicarbonate,  being  sparingly  soluble,  crystallizes  out. 
A  mixture  of  the  bicarbonate  with  some  substance  (the. 
compound  known  as  cream  of  tartar  is  generally  used) 
which  slowly  reacts  with  it,  liberating  carbon  dioxide,  is 
used  largely  in  baking.  The  carbon  dioxide  generated 
forces  its  way  through  the  dough,  thus  making  it  porous 
and  light. 

Sodium  nitrate  (Chili  saltpeter}  (NaNO3).  This  sub- 
stance is  found  in  nature  in  arid  regions  in  a  number  of 
places,  where  it  has  been  formed  apparently  by  the  decay 
of  organic  substances  in  the  presence  of  air  and  sodium 
salts.  The  largest  deposits  are  in  Chili,  and  most  of  the 
nitrate  of  commerce  comes  from  that  country.  Smaller 
deposits  occur  in  California  and  Nevada.  The  commercial 
salt  is  prepared  by  dissolving  the  crude  nitrate  in  water, 


286     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

allowing  the  insoluble  earthy  materials  to  settle,  and  evapo- 
rating the  clear  solution  so  obtained  to  crystallization.  The 
soluble  impurities  remain  for  the  most  part  in  the  mother 
liquors. 

Since  this  salt  is  the  only  nitrate  found  extensively  in 
nature,  it  is  the  material  from  which  other  nitrates  as  well 
as  nitric  acid  are  prepared.  It  is  used  in  enormous  quanti- 
ties in  the  manufacture  of  sulphuric  acid  and  potassium 
nitrate,  and  as  a  fertilizer. 

Sodium  phosphate  (Na2HPO4-  12  H2O).  Since  phosphoric 
acid  has  three  replaceable  hydrogen  atoms,  three  sodium 
phosphates  are  possible,  —  two  acid  salts  and  one  normal. 
All  three  can  be  made  without  difficulty,  but  disodium 
phosphate  is  the  only  one  which  is  largely  used,  and  is 
the  salt  which  is  commonly  called  sodium  phosphate.  It  is 
made  by  the  action  of  phosphoric  acid  on  sodium  carbonate : 

Na2CO3  +  H3PO4  =  Na2HPO4  +  CO2  +  H2O. 

It  is  interesting  as  being  one  of  the  few  phosphates  which 
are  soluble  in  water,  and  is  the  salt  commonly  used  when 
a  soluble  phosphate  is  needed. 

Normal  sodium  phosphate  (Na3PO4).  Although  this  is 
a  normal  salt  its  solution  has  a  strongly  alkaline  reaction. 
This  is  due  to  the  fact  that  the  salt  hydrolyzes  in  solution 
into  sodium  hydroxide  and  disodium  phosphate,  as  repre- 
sented in  the  equation 

Na3PO4  +  H2O  =  Na2HPO4  +  NaOH. 

Sodium  hydroxide  is  strongly  alkaline,  while  disodium  phos- 
phate is  nearly  neutral  in  reaction.  The  solution  as  a  whole 
is  therefore  alkaline.  The  salt  is  prepared  by  adding  a 
large  excess  of  sodium  hydroxide  to  a  solution  of  disodium 


THE  ALKALI  METALS  287 

phosphate  and  evaporating  to  crystallization.  The  excess 
of  the  sodium  hydroxide  reverses  the  reaction  of  hydrolysis 
and  the  normal  salt  crystallizes  out. 

Sodium  tetraborate  (borax)  (Na2B4O7-  10  H2O).  The 
properties  of  this  important  compound  have  been  discussed 
under  the  head  of  boron. 

POTASSIUM 

Occurrence  in  nature.  Potassium  is  a  constituent  of  many 
common  rocks  and  minerals,  and  is  therefore  a  rather 
abundant  element,  though  not  so  abundant  as  sodium. 
Feldspar,  which  occurs  both  by  itself  and  as  a  constituent 
of  granite,  contains  considerable  potassium.  The  element 
is  a  constituent  of  all  clay  and  of  mica  and  also  occurs  in 
very  large  deposits  at  Stassfurt,  Germany,  in  the  form  of 
the  chloride  and  sulphate,  associated  with  compounds  of 
sodium  and  magnesium.  In  small  quantities  it  is  found  as 
nitrate  and  in  many  other  forms. 

The  natural  decomposition  of  rocks  containing  potassium 
gives  rise  to  various  compounds  of  the  element  in  all  fertile 
soils.  Its  soluble  compounds  are  absorbed  by  growing 
plants  and  built  up  into  complex  vegetable  substances ; 
when  these  are  burned  the  potassium  remains  in  the  ash  in 
the  form  of  the  carbonate.  Crude  carbonate  obtained  from 
wood  ashes  was  formerly  the  chief  source  of  potassium 
compounds  ;  they  are  now  mostly  prepared  from  the  salts 
of  the  Stassfurt  deposits. 

Stassfurt  salts.  These  salts  form  very  extensive  deposits  in  middle 
and  north  Germany,  the  most  noted  locality  for  working  them  being 
at  Stassfurt.  The  deposits  are  very  thick  and  rest  upon  an  enormous 
layer  of  common  salt.  They  are  in  the  form  of  a  series  of  strata, 
each  consisting  largely  of  a  single  mineral  salt.  A  cross  section  of 


288     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


these  deposits  is  shown  in  Fig.  78.  While  these  strata  are  salts  from 
a  chemical  standpoint,  they  are  as  solid  and  hard  as  many  kinds  of 
stone,  and  are  mined  as  stone  or  coal  would  be.  Since  the  strata 
differ  in  general  appearance,  each  can  be  mined  separately,  and  the 
various  minerals  can  be  worked  up  by  methods  adapted  to  each  par- 
ticular case.  The  chief  minerals  of  commercial  importance  in  these 
deposits  are  the  following  : 


Sylvine 
Anhydrite 
Carnallite 
Kainite  . 
Polyhalite 
Kieserite  . 
Schonite  . 


6H2O. 


KC1. 

CaSO4. 

KC1  •  MgCl2 

K2SO4  •  MgS04  -  MgCl2  •  6  H2O. 

K2S04  •  MgSO4  •  2  CaSO4  •  2  H2O. 

MgS04-H20. 

K2S04-MgSO4-6H2O. 


Preparation  and  properties.   The  metal  is  prepared  by  the 
same  method  used  in  the  preparation  of  sodium.    In  most 

respects  it  is  very  simi- 
lar to  sodium,  the  chief 
difference  being  that  it 
is  even  more  energetic 
in  its  action  upon  other 
substances.  The  freshly 
cut,  bright  surface  in- 
stantly becomes  dim 
through  oxidation  by 
the  air.  It  decomposes 
water  very  vigorously, 
the  heat  of  reaction  be- 
ing sufficient  to  ignite  the 
hydrogen  evolved.  It  is 

Fir   78 

somewhat    lighter    than 
sodium  and  is  preserved  under  gasoline. 

Potassium  hydroxide  (caustic  potash)  (KOH).   Potassium 
hydroxide  is  prepared  by  methods  exactly  similar  to  those 


THE  ALKALI  METALS  289 

used  in  the  preparation  of  sodium  hydroxide,  which  com- 
pound it  closely  resembles  in  both  physical  and  chemical 
properties.  It  is  not  used  to  any  very  great  extent,  being 
replaced  by  the  cheaper  sodium  hydroxide. 

Action  of  the  halogen  elements  on  potassium  hydroxide. 
When  any  one  of  the  three  halogen  elements  —  chlorine, 
bromine,  and  iodine  —  is  added  to  a  solution  of  potassium 
hydroxide  a  reaction  takes  place,  the  nature  of  which 
depends  upon  the  conditions  of  the  experiment.  Thus, 
when  chlorine  is  passed  into  a  cold  dilute  solution  of 
potassium  hydroxide  the  reaction  expressed  by  the  follow- 
ing equation  takes  place  : 

(1)  2  KOH  +  2  Cl  =  KC1  +  KC10  +  H20. 

If  the  solution  of  hydroxide  is  concentrated  and  hot,  on  the 
other  hand,  the  potassium  hypochlorite  formed  according 
to  equation  .(i)  breaks  down  as  fast  as  formed  : 

(2)  3  KC1O  =  KC1O3  +  2  KC1. 

Equation  (i),  after  being  multiplied  by  3,  may  be  combined 
with  equation  (2),  giving  the  following : 

(3)  6  KOH  +  6  Cl  =  5  KC1  +  KC1O3  +  3  H2O. 

This  represents  in  a  single  equation  the  action  of  chlorine 
on  hot,  concentrated  solutions  of  potassium  hydroxide. 
By  means  of  these  reactions  one  can  prepare  potassium 
chloride,  potassium  hypochlorite,  and  potassium  chlorate. 
By  substituting  bromine  or  iodine  for  chlorine  the  corre- 
sponding compounds  of  these  elements  are  obtained.  Some 
of  these  compounds  can  be  obtained  in  cheaper  ways. 

If  the  halogen  element  is  added  to  a  solution  of  sodium 
hydroxide  or  calcium  hydroxide,  the  reaction  which  takes 
place  is  exactly  similar  to  that  which  takes  place  with 


290     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

potassium  hydroxide.  It  is  possible,  therefore,  to  prepare 
in  this  way  the  sodium  and  calcium  compounds  correspond- 
ing to  the  potassium  compounds  given  above. 

Potassium  chloride  (KC1).  This  salt  occurs  in  nature  in 
sea  water,  in  the  mineral  sylvine,  and,  combined  with  mag- 
nesium chloride,  as  carnallite  (KC1  •  MgCl2  •  6  H2O).  It  is 
prepared  from  carnallite  by  saturating  boiling  water  with 
the  mineral  and  allowing  the  solution  to  cool.  The  mineral 
decomposes  while  in  solution,  and  the  potassium  chloride 
crystallizes  out  on  cooling,  while  the  very  soluble  mag- 
nesium chloride  remains  in  solution.  The  salt  is  very  simi- 
lar to  sodium  chloride  both  in  physical  and  chemical 
properties.  It  is  used  in  the  preparation  of  nearly  all 
other  potassium  salts,  and,  together  with  potassium  sul- 
phate, is  used  as  a  fertilizer. 

Potassium  bromide  (KBr).  When  bromine  is  added  to  a 
hot  concentrated  solution  of  potassium  hydroxide  there  is 
formed  a  mixture  of  potassium  bromide  and  potassium 
bromate  in  accordance  with  the  reactions  already  discussed. 
There  is  no  special  use  for  the  bromate,  so  the  solution  is 
evaporated  to  dryness,  and  the  residue,  consisting  of  a  mix- 
ture of  the  bromate  and  bromide,  is  strongly  heated.  This 
changes  the  bromate  to  bromide,  as  follows  : 

KBrO3  =  KBr  +  3  O. 

The  bromide  is  then  crystallized  from  water,  forming 
large  colorless  crystals.  It  is  used  in  medicine  and  in 
photography. 

Potassium  iodide  (KI).  Potassium  iodide  may  be  made 
by  exactly  the  same  method  as  has  just  been  described  for 
the  bromide,  substituting  iodine  for  bromine.  It  is  more 
frequently  made  as  follows.  Iron  filings  are  treated  with 


THE  ALKALI  METALS  29  1 

iodine,  forming  the  compound  Fe3I8  ;  on  boiling  this  sub- 
stance with  potassium  carbonate  the  reaction  represented 
in  the  following  equation  occurs  : 

Fe3I8  +  4  K2CO3  =  Fe3O4  +  8  KI  +  4  CO2. 

Potassium  iodide  finds  its  chief  use  in  medicine. 

Potassium  chlorate  (KC1O3).  This  salt,  as  has  just  been 
explained,  can  be  made  by  the  action  of  chlorine  on  strong 
potassium  hydroxide  solutions.  The  chief  use  of  potassium 
chlorate  is  as  an  oxidizing  agent  in  the  manufacture  of 
matches,  fireworks,  and  explosives  ;  it  is  also  used  in  the 
preparation  of  oxygen  and  in  medicine. 

Commercial  preparation.  By  referring  to  the  reaction  between 
chlorine  and  hot  concentrated  solutions  of  potassium  hydroxide,  it 
will  be  seen  that  only  one  molecule  of  potassium  chlorate  is  formed 
from  six  molecules  of  potassium  hydroxide.  Partly  because  of  this 
poor  yield  and  partly  because  the  potassium  hydroxide  is  rather 
expensive,  this  process  is  not  an  economical  one  for  the  preparation 
of  potassium  chlorate.  The  commercial  method  is  the  following. 
Chlorine  is  passed  into  hot  solutions  of  calcium  hydroxide,  a  com- 
pound which  is  very  cheap.  The  resulting  calcium  chloride  and 
chlorate  are  both  very  soluble.  To  the  solution  of  these  salts  potas- 
sium chloride  is  added,  and  as  the  solution  cools  the  sparingly  soluble 
potassium  chlorate  crystallizes  out  : 


Ca(C103)2  +  2KC\  =  2  KC103  +  CaCl2. 
Electro-chemical  processes  are  also  used. 

Potassium  nitrate  (saltpeter]  (KNO3).  This  salt  was  for- 
merly made  by  allowing  animal  refuse  to  decompose  in  the 
open  air  in  the  presence  of  wood  ashes  or  earthy  materials 
containing  potassium.  Under  these  conditions  the  nitrogen 
in  the  organic  matter  is  in  part  converted  into  potassium 
nitrate,  which  was  obtained  by  extracting  the  mass  with 
water  and  evaporating  to  crystallization.  This  crude  and 


292     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

slow  process  is  now  almost  entirely  replaced  by  a  manu- 
facturing process  in  which  the  potassium  salt  is  made  from 
Chili  saltpeter : 

NaNO3  +  KC1  =  NaCl  +  KNO3. 

This  process  has  been  made  possible  by  the  discovery  of 
the  Chili  niter  beds  and  the  potassium  chloride  of  the 
Stassfurt  deposits. 

The  reaction  depends  for  its  success  upon  the  apparently  insig- 
nificant fact  that  sodium  chloride  is  almost  equally  soluble  in  cold 
and  hot  water.  All  four  factors  in  the  equation  are  rather  soluble 
in  cold  water,  but  in  hot  water  sodium  chloride  is  far  less  soluble 
than* the  other  three.  When  hot  saturated  solutions  of  sodium 
nitrate  and  potassium  chloride  are  brought  together,  sodium  chloride 
precipitates  and  can  be  filtered  off,  leaving  potassium  nitrate  in  solution, 
together  with  some  sodium  chloride.  On  cooling,  potassium  nitrate 
crystallizes, out,  leaving  small  amounts  of  the  other  salts  in  solution. 

Potassium  nitrate  is  a  colorless  salt  which  forms  very 
large  crystals.  It  is  stable  in  the  air,  and  when  heated  is  a 
good  oxidizing  agent,  giving  up  oxygen  quite  readily.  Its 
chief  use  is  in  the  manufacture  of  gunpowder. 

Gunpowder.  The  object  sought  for  in  the  preparation  of  gun- 
powder is  to  secure  a  solid  substance  which  will  remain  unchanged 
under  ordinary  conditions,  but  which  will  explode  readily  when 
ignited,  evolving  a  large  volume  of  gas.  When  a  mixture  of  carbon 
and  potassium  nitrate  is  ignited  a  great  deal  of  gas  is  formed,  as 
will  be  seen  from  the  equation 

2  KNO3  +  3  C  =  CO2  +  CO  +  N2  +  K2CO3. 

By  adding  sulphur  to  the  mixture  the  volume  of  gas  formed  in  the 
explosion  is  considerably  increased  : 

2  KNO3  +  3  C  +  S  -  3  CO2  +  N2  +  K2S. 

Gunpowder  is  simply  a  mechanical  mixture  of  these  three  substances 
in  the  proportion  required  for  the  above  reaction.  While  the  equa- 
tion represents  the  principal  reaction,  other  reactions  also  take  place. 


THE  ALKALI  METALS 


293 


The  gases  formed  in  the  explosion,  when  measured  under  standard 
conditions,  occupy  about  two  hundred  and  eighty  times  the  volume  of 
the  original  powder.  Potassium  sulphide  (K2S)  is  a  solid  substance, 
and  it  is  largely  due  to  it  that  gunpowder  gives  off  smoke  and  soot 
when  it  explodes.  Smokeless  powder  consists  of  organic  substances 
which,  on  explosion,  give  only  colorless  gases,  and  hence  produce 
no  smoke.  Sodium  nitrate  is  cheaper  than  potassium  nitrate,  but 
it  is  not  adapted  to  the  manufacture  of  the  best  grades  of  powder, 
since  it  is  somewhat  deliquescent  and  does  not  give  up  its  oxygen  so 
readily  as  does  potassium  nitrate.  It  is  used,  however,  in  the  cheaper 
grades  of  powder,  such  as  are  employed  for  blasting. 

Potassium  cyanide  (KCN).  When  animal  matter  con- 
taining nitrogen  is  heated  with  iron  and  potassium  carbon- 
ate, complicated  changes  occur  which  result  in  the  formation 
of  a  substance  commonly  called  yellow  prussiate  of  potash, 
which  has  the  formula  K4FeC6N6.  When  this  substance  is 
heated  with  potassium,  potassium  cyanide  is  formed  : 

K4FeC6N6  +  2  K  =  6  KCN  +  Fe. 

Since  sodium  is  much  cheaper  than  potassium  it  is  often 
used  in  place  of  it : 

K4FeC6N6  +  2  Na  =  4  KCN  +  2  NaCN  +  Fe. 
The  mixture  of  cyanides  so  resulting  serves  most  of  the 
purposes  of  the  pure  salt.  It  is  used  very  extensively  in 
several  metallurgical  processes,  particularly  in  the  extrac- 
tion of  gold.  Potassium  cyanide  is  a  white  solid  character- 
ized by  its  poisonous  properties,  and  must  be  used  with 
extreme  caution. 

Potassium  carbonate  (potash]  (K2CO3).  This  compound 
occurs  in  wood  ashes  in  small  quantities.  It  cannot  be 
prepared  by  the  Solvay  process,  since  the  acid  carbonate 
is  quite  soluble  in  water,  but  is  made  by  the  Le  Blanc 
process.  Its  chief  use  is  in  the  manufacture  of  other 
potassium  salts. 


294     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Other  salts  of  potassium.  Among  the  other  salts  of 
potassium  frequently  met  with  are  the  sulphate  (K2SO4), 
the  acid  carbonate  (KHCO3),  the  acid  sulphate  (KHSO4), 
and  the  acid  sulphite  (KHSO3).  These  are  all  white  solids. 

LITHIUM,  RUBIDIUM,  CESIUM 

Of  the  three  remaining  elements  of  the  family  —  lithium,  rubidium, 
and  caesium  —  lithium  is  by  far  the  most  common,  the  other  two  being 
very  rare.  Lithium  chloride  and  carbonate  are  not  infrequently  found 
in  natural  mineral  waters,  and  as  these  substances  are  supposed  to 
increase  the  medicinal  value  of  the  water,  they  are  very  often  added 
to  artificial  mineral  waters  in  small  quantities. 

COMPOUNDS  OF  AMMONIUM 

General.  As  explained  in  a  previous  chapter,  when 
ammonia  is  passed  into  water  the  two  compounds  combine 
to  form  the  base  NH4OH,  known  as  ammonium  hydroxide. 
When  this  base  is  neutralized  with  acids  there  are  formed 
the  corresponding  salts,  known  as  the ,  ammonium  salts. 
Since  the  ammonium  group  is  univalent,  ammonium  salts 
resemble  those  of  the  alkali  metals  in  formulas ;  they 
also  resemble  the  latter  salts  very  much  in  their  chemical 
properties,  and  may  be  conveniently  described  in  connec- 
tion with  them.  Among  the  ammonium  salts  the  chloride, 
sulphate,  carbonate,  and  sulphide  are  the  most  familiar. 

Ammonium  chloride  (sal  ammoniac]  (NH4C1).  This 
substance  is  obtained  by  neutralizing  ammonium  hydroxide 
with  hydrochloric  acid.  It  is  a  colorless  substance  crystal- 
lizing in  fine  needles,  and,  like  most  ammonium  salts,  is 
very  soluble  in  water.  When  placed  in  a  tube  and  heated 
strongly  it  decomposes  into  hydrochloric  acid  and  ammonia. 
When  these  gases  reach  a  cooler  portion  of  the  tube  they 


THE  ALKALI  METALS 


295 


at  once  recombine,  and  the  resulting  ammonium  chloride  is 
deposited  on  the  sides  of  the  tube.  In  this  way  the  salt 
can  be  separated  from  non-volatile  impurities.  Ammonium 
chloride  is  sometimes  used  in  preparation  of  ammonia ;  it 
is  also  used  in  making  dry  batteries  and  in  the  laboratory 
as  a  chemical  reagent. 

Ammonium  sulphate  ((NH4)2SO4).  This  salt  resembles 
the  chloride  very  closely,  and,  being  cheaper,  is  used  in 
place  of  it  when  possible.  It  is  used  in  large  quantity  as 
a  fertilizer,  the  nitrogen  which  it  contains  being  a  very 
valuable  food  for  plants. 

Ammonium  carbonate  ((NH4)2CO3).  This  salt,  as  well 
as  the  acid  carbonate  (NH4HCO3),  is  used  as  a  chemical 
reagent.  They  are  colorless  solids,  freely  soluble  in  water. 
The  normal  carbonate  is  made  by  heating  ammonium 
chloride  with  powdered  limestone  (calcium  carbonate),  the 
ammonium  carbonate  being  obtained  as  a  sublimate  in 
compact  hard  masses  : 

2  NH4C1  +  CaCO3  =  (NH4)2CO3  +  CaCl2. 

The  salt  always  smells  of  ammonia,  since  it  slowly  decom- 
poses, as  shown  in  the  equation 

(NH4)2C03  =  NH4HC03  +  NH3. 

The  acid  carbonate,  or  bicarbonate,  is  prepared  by  saturat- 
ing a  solution  of  ammonium  hydroxide  with  carbon  dioxide  : 

NH4OH  +  C02  =  NH4HC03. 

It  is  a  well-crystallized  stable  substance. 

Ammonium  sulphide  ((NH4)2S).  Ammonium  sulphide 
is  prepared  by  the  action  of  hydrosulphuric  acid  upon 
ammonium  hydroxide  : 

2  NH4OH  +  H2S  =  (NH4)2S  +  2  H2O. 


296     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

If  the  action  is  allowed  to  continue  until  no  more  hydrosul- 
phuric  acid  is  absorbed,  the  product  is  the  acid  sulphide, 
sometimes  called  the  hydrosulphide  : 

NH4OH  +  H2S  =  NH4HS  +  H2O. 

If  equal  amounts  of  ammonium  hydroxide  and  ammonium 
acid  sulphide  are  brought  together,  the  normal  sulphide  is 

d  :    NH4OH  +  NH4HS  =  (NH4)2S  +  H2Q 

It  has  been  obtained  in  the  solid  state,  but  only  with  great 
difficulty.  As  used  in  the  laboratory  it  is  always  in  the 
form  of  a  solution.  It  is  much  used  in  the  process  of 
chemical  analysis  because  it  is  a  soluble  sulphide  and 
easily  prepared.  On  exposure  to  the  air  ammonium  sul- 
phide slowly  decomposes,  being  converted  into  ammonia, 
water,  and  sulphur : 

(NH4)2S  +  O  =  2  NH3  +  H2O  +  S. 

As  fast  as  the  sulphur  is  liberated  it  combines  with  the 
unchanged  sulphide  to  form  several  different  ammonium 
sulphides  in  which  there  are  from  two  to  five  sulphur 
atoms  in  the  molecule,  thus :  (NH4)2S2,(NH4)2S3,  (NH4)2S5. 
These  sulphides  in  turn  decompose  by  further  action  of 
oxygen,  so  that  the  final  products  of  the  reaction  are  those 
given  in  the  equation.  A  solution  of  these  compounds  is 
yellow  and  is  sometimes  called  yellow  ammonium  sulphide. 

FLAME  REACTION  —  SPECTROSCOPE 

When  compounds  of  either  sodium  or  potassium  are  brought  into 
the  non-luminous  flame  of  a  Bunsen  burner  the  flame  becomes  colored. 
Sodium  compounds  color  it  intensely  yellow,  while  those  of  potassium 
color  it  pale  violet.  When  only  one  of  these  elements  is  present  it  is 


THE  ALKALI  METALS 


297 


easy  to  identify  it  by  this  simple  test,  but  when  both  are  present  the 
intense  color  of  the  sodium  flame  entirely  conceals  the  pale  tint  charac- 
teristic of  potassium  compounds. 

It  is  possible  to  detect  the  potassium  flame  in  such  cases,  however, 
in  the  following  way.  When  light  is  allowed  to  shine  through  a  very 
small  hole  or  slit  in  some  kind  of  a  screen,  such  as  a  piece  of  metal, 
upon  a  triangular  prism  of  glass,  the  light  is  bent  or  refracted  out  of 
its  course  instead  of  passing  straight  through  the  glass.  It  thus  comes 
out  of  the  prism  at  some  angle  to  the  line  at  which  it  entered.  Yellow 
light  is  bent  more  than  red,  and  violet  more  than  yellow.  When  light 
made  up  of  the  yellow  of  sodium  and  the  violet  of  potassium  shines 
through  a  slit  upon  such  a  prism,  the  yellow  and  the  violet  lights  come 


FIG.  79 

out  at  somewhat  different  angles,  and  so  two  colored  lines  of  light  —  a 
yellow  line  and  a  violet  line  —  are  seen  on  loojdng  into  the  prism  in 
the  proper  direction.  The  instrument  used  for  separating  the  rays  of 
light  in  this  way  is  called  a  spectroscope  (Fig.  79).  The  material  to 
be  tested  is  placed  on  a  platinum  wire  and  held  in  the  colorless  Bun- 
sen  flame.  The  resulting  light  passes  through  the  slit  in  the  end  of 
tube  J3,  and  then  through  B  to  the  prism.  The  resulting  lines  of  light 
are  seen  by  looking  into  the  tube  A,  which  contains  a  magnifying 
lens.  Most  elements  give  more  than  one  image  of  the  slit,  each  having 
a  different  color,  and  the  series  of  colored  lines  due  to  an  element  is 
called  its  spectrum. 


298     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

The  spectra  of  the  known  elements  have  been  carefully  studied, 
and  any  element  which  imparts  a  characteristic  color  to  a  flame,  or 
has  a  spectrum  of  its  own,  can  be  identified  even  when  other  elements 
are  present.  Through  the  spectroscopic  examination  of  certain  min- 
erals a  number  of  elements  have  been  discovered  by  the  observation 
of  lines  which  did  not  belong  to  any  known  element.  A  study  of  the 
substance  then  brought  to  light  the  new  element.  Rubidium  and 
caesium  were  discovered  in  this  way,  rubidium  having  bright  red  lines 
and  caesium  a  very  intense  blue  line.  Lithium  colors  the  flame  deep 
red,  and  has  a  bright  red  line  in  its  spectrum. 

EXERCISES 

1.  What  is  an  alkali  ?    Can  a  metal  itself  be  an  alkali  ? 

2.  Write   equations  showing   how  the   following    changes  may 
be  brought   about,  giving   the  general    principle  involved  in  each 
change:  NaCl — ^  Na2SO3,  Na2SO3  — ^  NaCl,  NaCl  — ^  NaBr, 
Na2SO4 — ^NaNO3,  NaNO3 — ^  NaHCO3. 

3.  What  carbonates  are  soluble  ? 

4.  State  the  conditions  under   which  the  reaction  represented 
by  the  following  equation  can  be  made  to  go  in  either  direction  : 

Na2CO3  -f  H2O  +  CO2  ^  2  NaHCO3. 

5.  Account  for  the  fact  that  solutions  of  sodium  carbonate  and 
potassium  carbonate  are  alkaline. 

6.  What  non-metallic  element  is  obtained  from  the  deposits  of 
Chili  saltpeter  ? 

7.  Supposing  concentrated  hydrochloric  acid  (den.  =  1.2)  to  be 
worth  six  cents  a  pound,  what  is  the  value  of  the  acid  generated  in  the 
preparation  of  I  ton  of  sodium  carbonate  by  the  Le  Blanc  process  ? 

8.  What  weight  of  sodium  carbonate  crystals  will  I  kg.  of  the 
anhydrous  salt  yield? 

9.  Write  equations  for  the  preparation  of  potassium  hydroxide  by 
three  different  methods. 

10.  What  would  take  place  if  a  bit  of  potassium  hydroxide  were 
left  exposed  to  the  air? 

11.  Write  the  equations  for  the  reactions  between  sodium  hydrox- 
ide and  bromine ;  between  potassium  hydroxide  and  iodine. 

12.  Write  equations  for  the  preparation  of  potassium  sulphate  ;  of 
potassium  acid  carbonate,  • 


THE  ALKALI  METALS  299 

13.  What  weight  of  carnallite  would  be  necessary  in  the  prepara- 
tion of  i  ton  of  potassium  carbonate  ? 

14.  Write  the  equations  showing  how  ammonium  chloride,  ammo- 
nium sulphate,  ammonium  carbonate,  and  ammonium  nitrate  may  be 
prepared  from  ammonium  hydroxide. 

15.  Write  an  equation  to  represent  the  reaction  involved  in  the 
preparation  of  ammonia  from  ammonium  chloride. 

16.  What  substances  already  studied  are  prepared  from  the  fol- 
lowing compounds?  ammonium  chloride  ;  ammonium  nitrate  ;  ammon' 
ium  nitrite  ;  sodium  nitrate  ;  sodium  chloride. 

17.  How  could  you  prove  that  the  water  in  crystals  of  commoii 
salt  is  not  water  of  crystallization? 

18.  How  could  you  distinguish  between  potassium  chloride  and 
potassium  iodide?  between  sodium  chloride  and  ammonium  chloride? 
between  sodium  nitrate  and  potassium  nitrate  ? 


CHAPTER    XXIV 
THE  ALKALINE-EARTH  FAMILY 


MILLIGRAMS  SOL- 

UBLE IN  I  L. 

SYMBOL 

ATOMIC 
WEIGHT 

DENSITY 

OF  WATER  AT  18° 

CARBONATE 
DECOMPOSES 

SUL- 

HYDROX- 

PHATE 

IDE 

Calcium     . 

Ca 

40.1 

i-54 

2070.00 

1,670. 

At  dull  red  heat 

Strontium  . 

Sr 

87.6 

2.50 

170.00 

7,460. 

At  white  heat 

Barium  .     . 

Ba 

1374 

3-75 

2.29 

36,300. 

Scarcely  at  all 

The  family.  The  alkaline-earth  family  consists  of  the 
very  abundant  element  calcium  and  the  much  rarer  elements 
strontium  and  barium.  They  are  called  the  alkaline-earth 
metals  because  their  properties  are  between  those  of  the 
alkali  metals  and  the  earth  metals.  The  earth  metals  will 
be  discussed  in  a  later  chapter.  The  family  is  also  fre- 
quently called  the  calcium  family. 

1.  Occurrence.    These   elements    do   not  occur  free    in 
nature.    Their  most  abundant  compounds  are  the  carbon- 
ates and  sulphates  ;  calcium  also  occurs  in  large  quantities 
as  the  phosphate  and  silicate. 

2.  Preparation.    The    metals    were    first    prepared    by 
Davy  in    1808   by  electrolysis.    This   method    has   again 
come  into   use   in   recent   years.    Strontium    and   barium 
have  as  yet  been  obtained  only  in  small  quantities  and  in 
the  impure  state,  and  many  of  their  physical  properties, 

300 


THE  ALKALINE-EARTH  FAMILY  301 

such  as  their  densities  and  melting  points,  are  therefore 
imperfectly  known. 

3.  Properties.    The  three   metals   resemble   each   other 
very  closely.    They  are  silvery-white  in  color  and  are  about 
as  hard  as  lead.    Their  densities  increase  with  their  atomic 
weights,  as  is  shown  in  the  table  on  opposite  page.   Like  the 
alkali  metals  they  have  a  strong  affinity  for  oxygen,  tarnish- 
ing in  the  air  through  oxidation.    They  decompose  water  at 
ordinary  temperatures,  forming  hydroxides  and  liberating 
hydrogen.    When  ignited  in  the  air  they  burn  with  bril- 
liancy, forming  oxides  of  the  general  formula  MO.    These 
oxides  readily  combine  with  water,  according  to  the  equa- 
tion M0  +  H20  -  M  (OH)2. 

Each  of  the  elements  has  a  characteristic  spectrum,  and 
the  presence  of  the  metals  can  easily  be  detected  by  the 
spectroscope. 

4.  Compounds.    The  elements  are  divalent  in  almost  all 
of  their  compounds,  and  these  compounds  in  solution  give 
simple,  divalent,  colorless  ions.    The  corresponding  salts 
of  the  three  elements  are  very  similar  to  each  other  and 
show   a   regular  variation   in   properties   in   passing  from 
calcium  to  strontium  and  from  strontium  to  barium.    This 
is  seen  in  the  solubility  of    the  sulphate  and  hydroxide, 
and  in  the  ease  of  decomposition  of  ,the  carbonates,  as 
given  in  the  table.     Unlike  the  alkali  metals,  their  normal 
carbonates  and  phosphates  are  insoluble. in  water. 

CALCIUM 

Occurrence.  The  compounds  of  calcium  are  very  abundant 
in  nature,  so  that  the  total  amount  of  calcium  in  the  earth's 
crust  is  very  large.  A  great  many  different  compounds 


302     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

containing  the  element  are  known,  the  most  important  of 
which  are  the  following : 

Calcite  (marble)  .     ..    .     .     .     .  CaCO3. 

Phosphorite Ca3(PO4)2. 

Fluorspar CaF2. 

Wollastonite     ...    «..,...*.  CaSiO3. 

Gypsum CaSO4  •  2  H2O. 

Anhydrite    .     .     ,     .     .     .     .     .  CaSO4. 

Preparation.  Calcium  is  now  prepared  by  the  electroly- 
.  sis  of  the  melted  chloride,  the  metal  depositing  in  solid 
condition  on  the  cathode.  It  is  a  gray  metal,  considerably 
heavier  and  harder  than  sodium.  It  acts  upon  water, 
forming  calcium  hydroxide  and  hydrogen,  but  the  action 
does  not  evolve  sufficient  heat  to  melt  the  metal.  It 
promises  to  become  a  useful  substance,  though  no  com- 
mercial applications  for  it  have  as  yet  been  found. 

Calcium  oxide  (lime,  quicklime]  (CaO).  Lime  is  pre- 
pared by  strongly  heating  calcium  carbonate  (limestone)  in 
large  furnaces  called  kilns  : 

CaCO3  =  CaO  +  CO2. 

When  pure,  lime  is  a  white  amorphous  substance. 
Heated  intensely,  as  in  the  oxyhydrogen  flame,  it  gives  a 
brilliant  light  called  the  lime  light.  Although  it  is  a  very 
difficultly  fusible  substance,  yet  in  the  electric  furnace  it 
can  be  made  to  melt  and  even  boil.  Water  acts  upon  lime 
with  the  evolution  of  a  great  deal  of  heat,  —  hence  the 
name  quicklime,  or  live  lime,  —  the  process  being  called 
slaking.  The  equation  is 

CaO  +  H2O  =  Ca(OH)2. 

Lime  readily  absorbs  moisture  from  the  air,  and  is  used  to 
dry   moist   gases,   especially   ammonia,   which   cannot   be 


THE  ALKALINE-EARTH  FAMILY 


303 


dried  by  the  usual  desiccating  agents.  It  also  absorbs 
carbon  dioxide,  forming  the  carbonate 

CaO  +  CO2  =  CaCO3. 

Lime  exposed  to  air  is  therefore  gradually  converted  into 
hydroxide  and  carbonate,  and  will  no  longer  slake  with 
water.  It  is  then  said  to  be  air-slaked. 

Limekilns.  The  older  kiln,  still  in  common  use,  consists  of  a  large 
cylindrical  stack  in  which  the  limestone  is  loosely  packed.  A  fire  is 
built  at  the  base  of  the  stack,  and  when  the  burning  is  complete  it  is 
allowed  to  die  out  and  the  lime  is  removed  from  the  kiln.  The  newer 
kilns  are  constructed  as  shown  in  Fig.  80.  A  number  of  fire  boxes 
are  built  around  the  lower  part  of  the 
kiln,  one  of  which  is  shown  at  B.  The 
fire  is  built  on  the  grate  F  and  the  hot 
products  of  combustion  are  drawn  up 
through  the  stack,  decomposing  the  lime- 
stone. The  kiln  is  charged  at  C,  and 
sometimes  fuel  is  added  with  the  limestone 
to  cause  combustion  throughout  the  con- 
tents of  the  kiln.  The  burned  lime  is 
raked  out  through  openings  in  the  bottom 
of  the  stack,  one  of  which  is  shown  at  D. 
The  advantage  of  this  kind  of  a  kiln  over 
the  older  form  is  that  the  process  is  con- 
tinuous, limestone  being  charged  in  at 
the  top  as  fast  as  the  lime  is  removed 
at  the  bottom. 


Calcium  hydroxide  (slaked  lime) 
(Ca(OH)2).  Pure  calcium  hydrox- 
ide is  a  light  white  .powder.  It  is 
sparingly  soluble  in  water,  forming 
a  solution  called  limewatcr,  which  is  often  used  in  medi- 
cine as  a  mild  alkali.  Chemically,  calcium  hydroxide  is  a 
moderately  strong  base,  though  not  so  strong  as  sodium 
hydroxide.  Owing  to  its  cheapness  it  is  much  used  in  the 


FIG. 80 


304     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

industries  whenever  an  alkali  is  desired.  A  number  of 
its  uses  have  already  been  mentioned.  It  is  used  in  the 
preparation  of  ammonia,  bleaching  powder,  and  potassium 
hydroxide.  It  is  also  used  to  remove  carbon  dioxide  and 
sulphur  compounds  from  coal  gas,  to  remove  the  hair  from 
hides  in  the  tanneries  (this  recalls  the  caustic  or  corrosive 
properties  of  sodium  hydroxide),  and  for  making  mortar. 

Mortar  is  a  mixture  of  calcium  hydroxide  and  sand. 
When  it  is  exposed  to  the  air  or  spread  upon  porous  mate- 
rials moisture  is  removed  from  it  partly  by  absorption  in 
the  porous  materials  and  partly  by  evaporation,  and  the 
mortar  becomes  firm,  or  sets.  At  the  same  time  carbon 
dioxide  is  slowly  absorbed  from  the  air,  forming  hard  cal- 
cium carbonate : 

Ca(OH)2  -f  CO2  =  CaCO3  +  H2O. 

By  this  combined  action  the  mortar  becomes  very  hard 
and  adheres  firmly  to  the  surface  upon  which  it  is  spread. 
The  sand  serves  to  give  body  to  the  mortar  and  makes  it 
porous,  so  that  the  change  into  carbonate  can  take  place 
throughout  the  mass.  It  also  prevents  too  much  shrinkage. 
Cement.  When  limestone  to  which  clay  and  sand  have 
been  added  in  certain  proportions  is  burned  until  it  is 
partly  fused  (some  natural  marl  is  already  of  about  the 
right  composition),  and  the  clinker  so  produced  is  ground 
to  powder,  the  product  is  called  cement.  When  this 
material  is  moistened  it  sets  to  a  hard  stone -like  mass 
which  retains  its  hardness  even  when  exposed  to  the  con- 
tinued action  of  water.  It  can  be  used  for  under-water 
work,  such  as  bridge  piers,  where  mortar  would  quickly 
soften.  Several  varieties  of  cement  are  made,  the  best 
known  of  which  is  Portland  cement. 


THE  ALKALINE-EARTH  FAMILY  305 

Growing  importance  of  cement.  Cement  is  rapidly  coming  into  use 
for  a  great  variety  of  purposes.  It  is  often  used  in  place  of  mortar  in 
the  construction  of  brick  buildings.  Mixed  with  crushed  stone  and 
sand  it  forms  concrete  which  is  used  in  foundation  work.  It  is  also 
used  in  making  artificial  stone,  terra-cotta  trimmings  for  buildings, 
artificial  stone  walks  and  floors,  and  the  like.  It  is  being  used  more 
and  more  for  making  many  articles  which  were  formerly  made  of 
wood  or  stone,  and  the  entire  walls  of  buildings  are  sometimes  made 
of  cement  blocks  or  of  concrete. 

Calcium  carbonate  (CaCO3).  This  substance  is  found  in 
a  great  many  natural  forms  to  which  various  names  have 
been  given.  They  may  be  classified  under  three  heads  : 

1.  Amorphous    carbonate.    This    includes    those    forms 
which  are  not  markedly  crystalline.    Limestone  is  the  most 
familiar  of  these  and  is  a  grayish  rock  usually  found  in 
hard  stratified  masses.    Whole  mountain  ranges  are  some- 
times made  up  of  this  material.    It  is  always  impure,  usually 
containing    magnesium   carbonate,   clay,    silica,    iron    and 
aluminium  compounds,  and  frequently  fossil  remains.    Marl 
is  a  mixture  of  limestone  and  clay.    Pearls,  chalk,  coral, 
and  shells  are  largely  calcium  carbonate. 

2.  Hexagonal  carbonate.    Calcium  carbonate  crystallizes 
in  the  form  of  rhomb-shaped  crystals  which  belong  to  the 
hexagonal  system.    When  very  pure  and  transparent  the 
substance  is  called  Iceland  spar.    Calcite  is  a  similar  form, 
but  somewhat  opaque  or  clouded.    Mexican  onyx  is  a  mas- 
sive variety,  streaked  or  banded  with  colors  due  to  impuri- 
ties.    Marble  when    pure   is   made   up   of  minute   calcite 
crystals.    Stalactites  and  stalagmites  are  icicle-like  forms 
sometimes  found  in  caves. 

3.  Rhombic    carbonate.    Calcium    carbonate    sometimes 
crystallizes    in    needle-shaped    crystals    belonging    to    the 
rhombic  system.    This  is  the  unstable  form  and  tends  to 


i 


306     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

go   over  into  the  other  variety.     Aragonite  is  the  most 
familiar  example  of  this  form. 

Preparation  and  uses  of  calcium  carbonate.    In  the  labo- 
ratory pure  calcium  carbonate  can  be  prepared  by  treating 
a  soluble  calcium  salt  with  a  soluble  carbonate  : 
Na2CO3  4-  CaCl2  =  CaCO3  +  2  NaCl. 

When  prepared  in  this  way  it  is  a  soft  white  powder  often 
called  precipitated  chalk,  and  is  much  used  as  a  polishing 
powder.  It  is  insoluble  in  water,  but  dissolves  in  water 
saturated  with  carbon  dioxide,  owing  to  the  formation  of 
the  acid  calcium  carbonate  which  is  slightly  soluble : 
CaCO3  +  H2CO3  =  Ca(HCO3)2. 

The  natural  varieties  of  calcium  carbonate  find  many 
uses,  such  as  in  the  preparation  of  lime  and  carbon  diox- 
ide ;  in  metallurgical  operations,  especially  in  the  blast 
furnaces ;  in  the  manufacture  of  soda,  glass,  and  crayon 
(which,  in  addition  to  chalk,  usually  contains  clay  and  cal- 
cium sulphate) ;  for  building  stone  and  ballast  for  roads. 

Calcium  chloride  (CaCl2).  This  salt  occurs  in  consider- 
able quantity  in  sea  water.  It  is  obtained  as  a  by-product 
in  many  technical  processes,  as  in  the  Solvay  soda  process. 
When  crystallized  from  its  saturated  solutions  it  forms 
colorless  needles  of  the  composition  CaCl2-  6  H2O.  By 
evaporating  a  solution  to  dryness  and  heating  to  a  mod- 
erate temperature  calcium  chloride  is  obtained  anhydrous 
as  a  white  porous  mass.  In  this  condition  it  absorbs  water 
with  great  energy  and  is  a  valuable  drying  agent. 

Bleaching  powder  (CaOCl2).  When  chlorine  acts  upon  a 
solution  of  calcium  hydroxide  the  reaction  is  similar  to  that 
which  occurs  between  chlorine  and  potassium  hydroxide  : 

2  Ca(OH)2  +  4  Cl  -  CaCl2  +  Ca(ClO)2  +  2  H2O. 


THE  ALKALINE-EARTH  FAMILY  307 

If,  however,  chlorine  is  conducted  over  calcium  hydroxide 
in  the  form  of  a  dry  powder,  it  is  absorbed  and  a  sub- 
stance is  formed  which  appears  to  have  the  composition 
represented  in  the  formula  CaOCl2.  This  substance  is 
called  bleaching  powder,  or  hypochlorite  of  lime.  It  is 
probably  the  calcium  salt  of  both  hydrochloric  and  hypo- 
chlorous  acids,  so  that  its  structure  is  represented  by  the 


formula  Ca^,       In  solution  this  substance  acts  exactly 

like  a  mixture  of  calcium  cnWride  (CaCl2)  and  calcium 
hypochlorite  (Ca(ClO)2),  since  it  dissociates  to  form  the 
ions  Ca++,  Q-,  and  ClCr. 

Bleaching  powder  undergoes  a  number  of  reactions 
which  make  it  an  important  substance. 

i.  When  treated  with  an  acid  it  evolves  chlorine  : 

Ca\Cl°  +  H2S°4  =  CaSO4  +  HC1  +  HC1O, 


+  HC1O  =  H2O  +  2  Cl. 

This  reaction  can  be  employed  in  the  preparation  of  chlo- 
rine, or  the  nascent  chlorine  may  be  used  as  a  bleaching 
agent. 

2.  It  is  slowly  decomposed  by  the  carbon  dioxide  of  the 
air,  yielding  calcium  carbonate  and  chlorine  : 

CaOCl2  +  CO2  =  CaCOg  +  2  Cl. 

Owing  to  this  slow  action  the  substance  is  a  good  disin- 
fectant. 

3.  When  its   solution  is   boiled   the   substance   breaks 
down  into  calcium  chloride  and  chlorate  : 

6  CaOCl2  =  5  CaCl2  +  Ca(ClO3)2. 

This  reaction  is  used  in  the  preparation  of  potassium 
chlorate. 


308     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Calcium  fluoride  (fluorspar)  (CaF2).  Fluorspar  has  already 
been  mentioned  as  the  chief  natural  compound  of  fluorine. 
It  is  found  in  large  quantities  in  a  number  of  localities, 
and  is  often  crystallized  in  perfect  cubes  of  a  light  green 
or  amethyst  color.  It  can  be  melted  easily  in  a  furnace, 
and  is  sometimes  used  in  the  fused  condition  in  metallur- 
gical operations  to  protect  a  metal  from  the  action  of  the 
air  during  its  reduction.  It  is  used  as  the  chief  source  of 
fluorine  compounds,  especially  hydrofluoric  acid. 

Calcium  sulphate  (gyps^M)  (CaSO4  •  2  H2O).  This  abun- 
dant substance  occurs  in  very  perfectly  formed  crystals 
or  in  massive  deposits.  It  is  often  found  in  solution  in 
natural  waters  and  in  the  sea  water.  Salts  deposited  from 
sea  water  are  therefore  likely  to  contain  this  substance 
(see  Stassfurt  salts). 

It  is  very  sparingly  soluble  in  water,  and  is  thrown  down 
as  a  fine  white  precipitate  when  any  considerable  amounts 
of  a  calcium  salt  and  a  soluble  sulphate  (or  ^^ftiric  acid) 
are  brought  together  in  solution.  Its  chief  use  is  in  the 
manufacture  of  plaster  of  Paris  and  of  hollow  tiles  for  fire- 
proof walls.  Such  material  is  called  gypsite.  It  is  also 
used  as  a  fertilizer. 

Calcium  sulphate,  like  the  carbonate,  occurs  in  many 
forms  in  nature.  Gypsum  is  a  name  given  to  all  common 
varieties.  Granular  or  massive  specimens  are  called  alabas- 
ter, while  all  those  which  are  well  crystallized  are  called 
selenite.  Satin  spar  is  still  another  variety  often  seen  in 
mineral  collections. 

Plaster  of  Paris.  When  gypsum  is  heated  to  about  115° 
it  loses  a  portion  of  its  water  of  crystallization  in  accord- 
ance with  the  equation 

2  (CaSO4  •  2  H2O)  =  2  CaSO4  -  H2O  +  3  H2O. 


THE  ALKALINE-EARTH  FAMILY  309 

The  product  is  a  fine  white  powder  called  plaster  of  Paris. 
On  being  moistened  it  again  takes  up  this  water,  and  in  so 
doing  first  forms  a  plastic  mass,  which  soon  becomes  very 
firm  and  hard  and  regains  its  crystalline  structure.  These 
properties  make  it  very  valuable  as  a  material  for  forming 
casts  and  stucco  work,  for  cementing  glass  to  metals,  and 
for  other  similar  purposes.  If  overheated  so  that  all  water 
is  driven  off,  the  process  of  taking  up  water  is  so  slow 
that  the  material  is  worthless.  Such  material  is  said  to  be 
dead  burned.  Plaster  of  Paris  is  very  extensively  used  as 
the  finishing  coat  for  plastered  walls. 

Hard  water.  Waters  containing  compounds  of  calcium 
and  magnesium  in  solution  are  called  hard  waters  because 
they  feel  harsh  to  the  touch.  The  hardness  of  water 
may  be  of  two  kinds,  —  (i)  temporary  hardness  and  (2) 
permanent  hardness. 

i .  Temporary  hardness.  We  have  seen  that  when  water 
charged  with  carbon  dioxide  comes  in  contact  with  lime- 
stone a  certain  amount  of  the  latter  dissolves,  owing  to 
the  formation  of  the  soluble  acid  carbonate  of  calcium. 
The  hardness  of  such  waters  is  said  to  be  temporary,  since 
it  may  be  removed  by  boiling.  The  heat  changes  the  acid 
carbonate  into  the  insoluble  normal  carbonate  which  then 
precipitates,  rendering  the  water  soft : 

Ca(HCO3)2  =  CaCO3  +  H2O  4-  CO2. 

Such  waters  may  also  be  softened  by  the  addition  of  suffi- 
cient lime  or  calcium  hydroxide  to  convert  the  acid  car- 
bonate of  calcium  into  the  normal  carbonate.  The  equation 
representing  the  reaction  is 

Ca(HC03)2  +  Ca(OH)2  =  2  CaCO3  +  2  H2O. 


310     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

2.  Permanent  hardness.  Tfye  hardness  of  water  may 
also  be  due  to  the  presence  of  calcium  and  magnesium 
sulphates  or  chlorides.  Boiling  the  water  does  not  affect 
these  salts  ;  hence  such  waters  are  said  to  have  permanent 
hardness.  They  may  be  softened,  however,  by  the  addition 
of  sodium  carbonate,  which  precipitates  the  calcium  and 
magnesium  as  insoluble  carbonates  : 

CaSO4  +  Na2CO3  =  CaCO3  +  Na2SO4. 
This  process  is  sometimes  called  "breaking"  the  water. 

Commercial  methods  for  softening  water.  The  average  water  of  a 
city  supply  contains  not  only  the  acid  carbonates  of  calcium  and 
magnesium  but  also  the  sulphates  and  chlorides  of  these  metals, 
together  with  other  salts  in  smaller  quantities.  Such  waters  are 
softened  on  a  commercial  scale  by  the  addition  of  the  proper  quan- 
tities of  calcium  hydroxide  and  sodium  carbonate.  The  calcium 
hydroxide  is  added  first  to  precipitate  all  the  acid  carbonates.  After 
a  short  time  the  sodium  carbonate  is  added  to  precipitate  the  other 
soluble  salts  of  calcium  and  magnesium,  together  with  any  excess  of 
calcium  hydroxide  which  may  have  been  added.  The  quantity  of 
calcium  hydroxide  and  sodium  carbonate  required  is  calculated  from 
a  chemical  analysis  of  the  water.  It  will  be  noticed  that  the  water 
softened  in  this  way  will  contain  sodium  sulphate  and  chloride,  but 
the  presence  of  these  salts  is  not  objectionable. 

Calcium  carbide  (CaC2).  This  substance  is  made  by 
heating  well-dried  coke  and  lime  in  an  electrical  furnace. 
The  equation  is 

CaO  +  3  C  =  CaC2  +  CO. 

The  pure  carbide  is  a  colorless,  transparent,  crystalline  sub- 
stance. In  contact  with  water  it  is  decomposed  with  the 
evolution  of  pure  acetylene  gas,  having  a  pleasant  ethereal 
odor.  The  commercial  article  is  a  dull  gray  porous  sub- 
stance which  contains  many  impurities.  The  acetylene  pre- 
pared from  this  substance  has  a  very  characteristic  odor 


THE  ALKALINE-EARTH  FAMILY 


due  to  impurities,  the  chief  of  these  being  phosphine.  It 
is  used  in  considerable  quantities  as  a  source  of  acetylene 
gas  for  illuminating  purposes. 

Technical  preparation.  Fig.  8 1  represents  a  recent  type  of  a  car- 
bide furnace.  The  base  of  the  furnace  is  provided  with  a  large  block 
of  carbon  A,  which  serves  as 
one  of  the  electrodes.  The 
other  electrodes  B,  several  in 
number,  are  arranged  hori- 
zontally at  some  distance 
above  this.  A  mixture  of  coal 
and  lime  is  fed  into  the 
furnace  through  the  trap  top  C, 
and  in  the  lower  part  of  the 
furnace  this  mixture  becomes 
intensely  heated,  forming 
liquid  carbide.  This  is  drawn 
off  through  the  taphole  D. 

The  carbon  monoxide 
formed  in  the  reaction  escapes 
through  the  pipes  E  and  is 
led  back  into  the  furnace.  The 
pipes  ^supply  air,  so  that  the 
monoxide  burns  as  it  reenters 
the  furnace  and  assists  in 

heating  the  charge.  The  carbon  dioxide  so  formed,  together  with 
the  nitrogen  entering  as  air,  escape  at  G.  An  alternating  current 
is  used. 

Calcium  phosphate  (Ca3(PO4)2).  This  important  substance 
occurs  abundantly  in  nature  as  a  constituent  of  apatite 
(3  Ca3(PO4)2  •  CaF2),  in  phosphate  rock,  and  as  the  chief 
mineral  constituent  of  bones.  Bone  ash  is  therefore  nearly 
pure  calcium  phosphate.  It  is  a  white  powder,  insoluble  in 
water,  although  it  readily  dissolves  in  acids,  being  decom- 
posed by  them  and  converted  into  soluble  acid  phosphates, 
as  explained  in  connection  with  the  acids  of  phosphorus. 


FIG. 81 


312     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

STRONTIUM 

Occurrence.  Strontium  occurs  sparingly  in  nature,  usu- 
ally as  strontianite  (SrCO3)  and  as  celestite  (SrSO4).  Both 
minerals  form  beautiful  colorless  crystals,  though  celestite 
is  sometimes  colored  a  faint  blue.  Only  a  few  of  the  com- 
pounds of  strontium  have  any  commercial  applications. 

Strontium  hydroxide  (Sr(OH)2- 8  H2O).  The  method 
of  preparation  of  strontium  hydroxide  is  analogous  to  that 
of  calcium  hydroxide.  The  substance  has  the  property  of 
forming  an  insoluble  compound  with  sugar,  which  can 
easily  be  separated  again  into  its  constituents.  It  is  there- 
fore sometimes  used  in  the  sugar  refineries  to  extract  sugar 
from  impure  mother  liquors  from  which  the  sugar  will 
not  crystallize. 

Strontium  nitrate  (Sr(NO3)2  •  4  H2O).  This  salt  is  pre- 
pared by  treating  the  native  carbonate  with  nitric  acid. 
When  ignited  with  combustible  materials  it  imparts  a  bril- 
liant crimson  color  to  the  flame,  and  because  of  this  prop- 
erty it  is  used  in  the  manufacture  of  red  lights. 

BARIUM 

Barium  is  somewhat  more  abundant  than  strontium, 
occurring  in  nature  largely  as  barytes,  or  heavy  spar 
(BaSO4),  and  witherite  (BaCO3).  Like  strontium,  it  closely 
resembles  calcium  both  in  the  properties  of  the  metal  and 
in  the  compounds  which  it  forms. 

Oxides  of  barium.  Barium  oxide  (BaO)  can  be  obtained 
by  strongly  heating  the  nitrate  : 

Ba(NO3)2  =  BaO  +  2  NO2  -f  O. 

Heated  to  a  low  red  heat  in  the  air,  the  oxide  combines  with 
oxygen,  forming  the  peroxide  (BaO2).     If  the  temperature 


THE  ALKALINE-EARTH  FAMILY  313 

is  raised  still  higher,  or  the  pressure  is  reduced,  oxygen  is 
given  off  and  the  oxide  is  once  more  formed.    The  reaction 


is  reversible  and  has  been  used  as  a  means  of  separating 
oxygen  from  the  air.  Treated  with  acids,  barium  peroxide 
yields  hydrogen  peroxide  : 

BaO2  +  2.  HC1  =  BaCl2  +  H2O2. 

Barium  chloride  (BaCl2  •  2  H2O).  Barium  chloride  is  a 
white  well-crystallized  substance  which  is  easily  prepared 
from  the  native  carbonate.  It  is  largely  used  in  the  labora- 
tory as  a  reagent  to  detect  the  presence  of  sulphuric  acid 
or  soluble  sulphates. 

Barium  sulphate  (barytes)  (BaSO4).  Barium  sulphate 
occurs  in  nature  in  the  form  of  heavy  white  crystals.  It 
is  precipitated  as  a  crystalline  powder  when  a  barium  salt 
is  added  to  a  solution  of  a  sulphate  or  sulphuric  acid  : 

BaCl2  +  H2SO4  =  BaSQ4  -f  2  HC1. 

This  precipitate  is  used,  as  are  also  the  finely  ground 
native  sulphate  and  carbonate,  as  a  pigment  in  paints. 
On  account  of  its  low  cost  it  is  sometimes  used  as  an 
adulterant  of  white  lead,  which  is  also  a  heavy  white 
substance. 

Barium  compounds  color  the  flame  green,  and  the 
nitrate  (Ba(NO3)2)  is  used  in  the  manufacture  of  green 
lights.  Soluble  barium  compounds  are  poisonous. 

RADIUM 

Historical.  In  1896  the  French  scientist  Becquerel  observed  that 
the  mineral  pitchblende  possesses  certain  remarkable  properties.  It 
affects  photographic  plates  even  in  complete  darkness,  and  discharges 


314     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

a  gold-leaf  electroscope  when  brought  close  to  it.  In  1898  Madam 
Curie  made  a  careful  study  of  pitchblende  to  see  if  these  properties 
belong  to  it  or  to  some  unknown  substance  contained  in  it.  She 
succeeded  in  extracting  from  it  a  very  small  quantity  of  a  substance 
containing  a  new  element  which  she  named  radium. 

Radium  itself  has  not  been  prepared,  but  its  compounds  so  closely 
resemble  those  of  barium  that  it  is  very  difficult  to  separate  the  two. 
Its  atomic  weight  is  about  two  hundred  and  twenty-five,  and  this 
value  places  it  in  the  same  family  with  barium. 

Properties.  Compounds  of  radium  affect  a  photographic  plate  or 
electroscope  even  through  layers  of  paper  or  sheets  of  metal.  They 
also  bring  about  chemical  changes  in  substances  placed  near  them. 
Investigation  of  these  strange  properties  has  suggested  that  the 
radium  atoms  are  unstable  and  undergo  a  decomposition.  As  a 
result  of  this  decomposition  very  minute  bodies,  to  which  the  name 
corpuscles  has  been  given,  are  projected  from  the  radium  atom  with 
exceedingly  great  velocity.  It  is  to  these  corpuscles  that  the  strange 
properties  of  radium  are  due.  It  seems  probable  that  the  gas  helium 
is  in  some  way  formed  during  the  decomposition  of  radium. 

Two  or  three  other  elements,  particularly  uranium  and  thorium, 
have  been  found  to  possess  many  of  the  properties  of  radium  in  smaller 
degree. 

Radium  and  the  atomic  theory.  If  these  views  in  regard  to 
radium  should  prove  to  be  well  founded,  it  will  be  necessary  to 
modify  in  some  respects  the  conception  of  the  atom  as  developed  in 
a  former  chapter.  The  atom  would  have  to  be  regarded  as  a  com- 
pound unit  made  up  of  several  parts.  In  a  few  cases,  as  in  radium 
and  uranium,  it  would  appear  that  this  unit  is  unstable  and  under- 
goes transformation  into  more  stable  combinations.  This  modifica- 
tion would  not,  in  any  essential  way,  be  at  variance  with  the  atomic 
theory  as  propounded  by  Dalton. 


EXERCISES 

1.  What  properties  have   the  alkaline-earth  metals  in  common 
with  the  alkali  metals?    In  what  respects  do  they  differ? 

2.  Write  the  equation  for  the  reaction  between  calcium  carbide 
and  water. 

3.  For  what  is  calcium  chlorate  used  ? 


THE  ALKALINE-EARTH  FAMILY  315 

4.  Could  limestone  be  completely  decomposed  if  heated  in  a 
closed  vessel  ? 

5.  Caves  often  occur  in  limestone.    Account  for  their  formation. 

6.  What  is  the  significance  of   the   term   fluorspar?    (Consult 
dictionary.) 

7.  Could  calcium  chloride  be  used  in  place  of  barium  chloride  in 
testing  for  sulphates  ? 

8.  What  weight  of  water  is  necessary  to  slake  the  lime  obtained 
from  i  ton  of  pure  calcium  carbonate  ? 

9.  What  weight  of  gypsum  is  necessary  in  the  preparation  of 
i  ton  of  plaster  of  Paris  ? 

10.  Write  equations  to  represent  the  reactions  involved  in  the 
preparation    of    strontium    hydroxide    and    strontium    nitrate    from 
strontianite. 

11.  Write  equations  to  represent  the  reactions   involved  in  the 
preparation  of  barium  chloride  from  heavy  spar. 

12.  Could  barium  hydroxide  be  used  in  place  of  calcium  hydrox- 
ide in  testing  for  carbon  dioxide? 


CHAPTER    XXV 
THE   MAGNESIUM  FAMILY 


SYMBOL 

ATOMIC 
WEIGHT 

DENSITY 

MELTING 
POINT 

BOILING 
POINT 

OXIDE 

Magnesium  . 
Zinc     .     .    ,     .   .-. 

Mg 
Zn 

24.36 

6  C  A 

!-75 

7  OO 

750° 
420° 

920° 
CKO° 

MgO 
ZnO 

Cadmium      . 

Cd 

II2.4 

8.67 

320° 

yyj 

778° 

CdO 

The  family.  In  the  magnesium  family  are  included  the 
four  elements :  magnesium,  zinc,  cadmium,  and  mercury. 
Between  the  first  three  of  these  metals  there  is  a  close 
family  resemblance,  such  as  has  been  traced  between  the 
members  of  the  two  preceding  families.  Mercury  in  some 
respects  is  more  similar  to  copper  and  will  be  studied  in 
connection  with  that  metal. 

1.  Properties.    When  heated  to  a  high  temperature  in 
the  air  each  of  these  metals  combines  with  oxygen  to  form 
an  oxide  of  the  general  formula  MO,  in  which  M  repre- 
sents the   metal.     Magnesium   decomposes   boiling  water 
slowly,   while   zinc   and    cadmium    have   but   little   action 
on  it. 

2.  Compounds.    The  members  of  this  group  are  divalent 
in  nearly  all  their  compounds,  so   that   the   formulas   of 
their  salts   resemble   those   of   the   alkaline-earth   metals. 
Like  the  alkaline-earth  metals,  their  carbonates  and  phos- 
phates are  insoluble  in  water.  Their  sulphates,  however,  are 
readily  soluble.    Unlike  both  the  alkali  and  alkaline-earth 

316 


THE  MAGNESIUM  FAMILY  317 

metals,  their  hydroxides  are  nearly  insoluble  in  water. 
Most  of  their  compounds  dissociate  in  such  a  way  as  to 
give  a  simple,  colorless,  metallic  ion. 


MAGNESIUM 

Occurrence.  Magnesium  is  a  very  abundant  element  in 
nature,  ranking  a  little  below  calcium  in  this  respect. 
Like  calcium,  it  is  a  constituent  of  many  rocks  and  also 
occurs  in  the  form  of  soluble  salts. 

Preparation.  The  metal  magnesium,  like  most  metals 
whose  oxides  are  difficult  to  reduce  with  carbon,  was 
formerly  prepared  by  heating  the  anhydrous  chloride  with 

MgCl2  +  2  Na  =  2  NaCl  +  Mg. 

It  is  now  made  by  electrolysis,  but  instead  of  using  as  the 
electrolyte  the  melted  anhydrous  chloride,  which  is  diffi- 
cult to  obtain,  the  natural  mineral  carnallite  is  used.  This 
is  melted  in  an  iron  pot  which  also  serves  as  the  cathode 
in  the  electrolysis.  A  rod  of  carbon  dipping  into  the 
melted  salt  serves  as  the  anode.  The  apparatus  is  very 
similar  to  the  one  employed  in  the  preparation  of  sodium. 
Properties.  Magnesium  is  a  rather  tough  silvery-white 
metal  of  small  density.  Air  does  not  act  rapidly  upon  it, 
but  a  thin  film  of  oxide  forms  upon  ita  surface,  dimming 
its  bright  luster.  The  common  acids  dissolve  it  with  the 
formation  of  the  corresponding  salts.  It  can  be  ignited 
readily  and  in  burning  liberates  much  heat  and  gives  a 
brilliant  white  light.  This  light  is  very  rich  in  the  rays 
which  affect  photographic  plates,  and  the  metal  in  the 
form  of  fine  powder  is  extensively  used  in  the  production 
of  flash  lights  and  for  white  lights  in  pyrotechnic  displays. 


3l8     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Magnesium  oxide  (magnesia]  (MgO).  Magnesium  oxide, 
sometimes  called  magnesia  or  magnesia  usta,  resembles 
lime  in  many  respects.  It  is  much  more  easily  formed 
than  lime  and  can  be  made  in  the  same  way, — by  igniting 
the  carbonate.  It  is  a  white  powder,  very  soft  and  light, 
and  is  unchanged  by  heat  even  at  very  high  temperatures. 
For  this  reason  it  is  used  in  the  manufacture  of  crucibles, 
for  lining  furnaces,  and  for  other  purposes  where  a  refrac- 
tory substance  is  needed.  It  combines  with  water  to  form 
magnesium  hydroxide,  but  much  more  slowly  and  with  the 
production  of  much  less  heat  than  in  the  case  of  calcium 
oxide. 

Magnesium  hydroxide  (Mg(OH)2).  The  hydroxide 
formed  in  this  way  is.  very  slightly  soluble  in  water,  but 
enough  dissolves  to  give  the  water  an  alkaline  reaction. 
Magnesium  hydroxide  is  therefore  a  fairly  strong  base. 
It  is  an  amorphous  white  substance.  Neither  magnesia 
nor  magnesium  salts  have  a  very  marked  effect  upon  the 
system ;  and  for  this  reason  magnesia  is  a  very  suitable 
antidote  for  poisoning  by  strong  acids,  since  any  excess 
introduced  into  the  system  will  have  no  injurious  effect. 

Magnesium  cement.  A  paste  of  magnesium  hydroxide  and  water 
slowly  absorbs  carbon  dioxide  from  the  air  and  becomes  very  hard. 
The  hardness  of  the  product  is  increased  by  the  presence  of  a  con- 
siderable amount  of  magnesium  chloride  in  the  paste.  The  hydrox- 
ide, with  or  without  the  chloride,  is  used  in  the  preparation  of  cements 
for  some  purposes. 

Magnesium  carbonate  (MgCO3).  Magnesium  carbonate 
is  a  very  abundant  mineral.  It  occurs  in  a  number  of 
localities  as  magnesite,  which  is  usually  amorphous,  but 
sometimes  forms  pure  crystals  resembling  calcite.  More 
commonly  it  is  found  associated  with  calcium  carbonate. 


THE  MAGNESIUM  FAMILY 


319 


The  mineral  dolomite  has  the  composition  CaCO3  •  MgCO3. 
Limestone  containing  smaller  amounts  of  magnesium  car- 
bonate is  known  as  dolomitic  limestone.  Dolomite  is  one 
of  the  most  common  rocks,  forming  whole  mountain 
masses.  It  is  harder  and  less  readily  attacked  by  acids 
than  limestone.  It'  is  valuable  as  a  building  stone  and  as 
ballast  for  roadbeds  and  foundations.  Like  calcium  car- 
bonate, magnesium  carbonate  is  insoluble  in  water,  though 
easily  dissolved  by  acids. 

Basic  carbonate  of  magnesium.  We  should  expect  to  find 
magnesium  carbonate  precipitated  when  a  soluble  magne- 
sium salt  and  a  soluble  carbonate  are  brought  together : 

Na2CO3  +  MgCl2  =  MgCO3  +  2  NaCl. 

Instead  of  this,  some  carbon  dioxide  escapes  and  the 
product  is  found  to  be  a  basic  carbonate.  The  most 
common  basic  carbonate  of  magnesium  has  the  formula 
4MgCO3  •  Mg(OH)2,  and  is  sometimes  called  magnesia  alba. 
This  compound  is  formed  by  the  partial  hydrolysis  of  the 
normal  carbonate  at  first  precipitated  : 

5  MgC03  +  2  H20  =  4  MgC03  •  Mg(OH)2  +  H2CO3. 

Magnesium  chloride  (MgCl2  •  6  H2O).  Magnesium  chloride 
is  found  in  many  natural  waters  and  in  many  salt  deposits 
(see  Stassfurt  salts).  It  is  obtained  as  a' by -product  in  the 
manufacture  of  potassium  chloride  from  carnallite.  As 
there  is  no  very  important  use  for  it,  large  quantities 
annually  go  to  waste.  When  heated  to  drive  off  the 
water  of  crystallization  the  chloride  is  decomposed  as 
shown  in  the  equation 

MgCl2  •  6  H2O  =  MgO  +  2  HC1  +  5  H2O. 


320     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Owing  to  the  abundance  of  magnesium  chloride,  this  reac- 
tion is  being  used  to  some  extent  in  the  preparation  of 
both  magnesium  oxide  and  hydrochloric  acid. 

Boiler  scale.  When  water  which  contains  certain  salts  in  solution 
is  evaporated  in  steam  boilers,  a  hard  insoluble  material  called  scale 
deposits  in  the  boiler.  The  formation  of  this  scale  may  be  due  to 
several  distinct  causes. 

1 .  To  the  deposit  of  calcium  sulphate.    This  salt,  while  sparingly 
soluble  in  cold  water,  is  almost  completely  insoluble  in  superheated 
water.    Consequently  it  is  precipitated  when  water  containing  it  is 
heated  in  a  boiler. 

2.  To  decomposition  of  acid  carbonates.    As  we  have  seen,  calcium 
and  magnesium  acid  carbonates  are  decomposed  on  heating,  forming 
insoluble  normal  carbonates: 

Ca(HCO3)2  =  CaCO3  +  H2O  +  CO2. 

3.  To  hydrolysis  of  magnesium  salts.    Magnesium  chloride,  and  to 
some  extent  magnesium  sulphate,  undergo  hydrolysis  when  super- 
heated in  solution,  and  the  magnesium  hydroxide,  being  sparingly 
soluble,  precipitates : 

MgCl2  +  2  H20  ^  Mg(OH)2  +  2  HC1. 

This  scale  adheres  tightly  to  the  boiler  in  compact  layers  and,  being 
a  non-conductor  of  heat,  causes  much  waste  of  fuel.  It  is  very  diffi- 
cult to  remove,  owing  to  its  hardness  and  resistance  to  reagents. 
Thick  scale  sometimes  cracks,  and  the  water  coming  in  contact  with 
the  overheated  iron  occasions  an  explosion.  Moreover,  the  acids  set 
free  in  the  hydrolysis  of  the  magnesium  salts  attack  the  iron  tubes 
and  rapidly  corrode  them.  These  causes  combine  to  make  the  forma- 
tion of  scale  a  matter  which  occasions  much  trouble  in  cases  where 
hard  water  is  used  in  steam  boilers.  Water  containing  such  salts 
should  be  softened,  therefore,  before  being  used  in  boilers. 

Magnesium  sulphate  (Epsom  salt}  (MgSO4  •  7  H2O).  Like 
the  chloride,  magnesium  sulphate  is  found  rather  commonly 
in  springs  and  in  salt  deposits.  A  very  large  deposit  of  the 
almost  pure  salt  has  been  found  in  Wyoming.  Its  name 


THE  MAGNESIUM  FAMILY  321 

was  given  to  it  because  of  its  abundant  occurrence  in  the 
waters  of  the  Epsom  springs  in  England. 

Magnesium  sulphate  has  many  uses  in  the  industries.  It 
is  used  to  a  small  extent  in  the  preparation  of  sodium  and 
potassium  sulphates,  as  a  coating  for  cotton  cloth,  in  the 
dye  industry,  in  tanning,  and  in  the  manufacture  of  paints 
and  laundry  soaps.  To  some  extent  it  is  used  in  medicine. 

Magnesium  silicates.  Many  silicates  containing  mag- 
nesium are  known  and  some  of  them  are  important  sub- 
stances. Serpentine,  asbestos,  talc,  and  meerschaum  are 
examples  of  such  substances. 

ZINC 

Occurrence.  Zinc  never  occurs  free  in  nature.  Its  com- 
pounds have  been  found  in  many  different  countries,  but 
it  is  not  a  constituent  of  common  rocks  and  minerals,  and 
its  occurrence  is  rather  local  and  confined  to  definite 
deposits  or  pockets.  It  occurs  chiefly  in  the  following  ores : 

Sphalerite  (zinc  blende)    ....  ZnS. 

Zincite ZnO. 

Smithsonite ZnCO3. 

Willemite Zn2SiO4. 

Franklinite ZnO  •  Fe2O3. 

One  fourth  of  the  world's  output  of  zinc  comes  from  the 
United  States,  Missouri  being  the  largest  producer. 

Metallurgy.  The  ores  employed  in  the  preparation  of 
zinc  are  chiefly  the  sulphide,  oxide,  and  carbonate.  They 
are  first  roasted  in  the  air,  by  which  process  they  are 
changed  into  oxide  : 

ZnCO3  =  ZnO  +  CO2, 
ZnS  +3O  =  ZnO  +  SO2. 


322      AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

The  oxide  is  then  mixed  with  coal  dust,  and  the  mixture  is 
heated  in  earthenware  muffles  or  retorts,  natural  ga$ 
being  used  as  fuel  in  many  cases.  The  oxide  is  reduced 
by  this  means  to  the  metallic  state,  and  the  zinc,  being 
volatile  at  the  high  temperature  reached,  distills  and  is  col- 
lected in  suitable  receivers.  At  first  the  zinc  collects  in  the 
form  of  fine  powder,  called  zinc  dust  or  flowers  of  zinc, 
recalling  the  formation  under  similar  conditions  of  flowers  of 
sulphur.  Later,  when  the  whole  apparatus  has  become  warm, 
the  zinc  condenses  to  a  liquid  in  the  receiver,  from  which  it 
is  drawn  off  into  molds.  Commercial  zinc  often  contains  a 
number  of  impurities,  especially  carbon,  arsenic,  and  iron. 

Physical  properties.  Pure  zinc  is  a  rather  heavy  bluish- 
white  metal  with  a  high  luster.  It  melts  at  about  420°,  and 
if  heated  much  above  this  temperature  in  the  air  takes  fire 
and  burns  with  a  very  bright  bluish  flame.  It  boils  at 
about  950°  and  can  therefore  be  purified  by  distillation. 

Many  of  the  physical  properties  of  zinc  are  much  influ- 
enced by  the  temperature  and  previous  treatment  of  the 
metal.  When  cast  into  ingots  from  the  liquid  state  it 
becomes  at  ordinary  temperatures  quite  hard,  brittle,  and 
highly  crystalline.  At  1 50°  it  is  malleable  and  can  be  rolled 
into  thin  sheets ;  at  higher  temperatures  it  again  becomes 
very  brittle.  When  once  rolled  into  sheets  it  retains  its 
softness  and  malleability  at  ordinary  temperatures.  When 
melted  and  poured  into  water  it  forms  thin  brittle  flakes, 
and  in  this  condition  is  called  granulated  or  mossy  zinc. 

Chemical  properties.  Zinc  is  tarnished  superficially  by 
moist  air,  but  beyond  this  is  not  affected  by  it.  It  does 
not  decompose  even  boiling  water.  When  the  metal  is 
quite  pure,  sulphuric  and  hydrochloric  acids  have  scarcely 
any  action  upon  it ;  when,  however,  it  contains  small 


THE  MAGNESIUM  FAMILY 


323 


amounts  of  other  metals  such  as  magnesium  or  arsenic,  or 
when  it  is  merely  in  contact  with  metallic  platinum,  brisk 
action  takes  place  and  hydrogen  is  evolved.  For  this  rea- 
son, when  pure  zinc  is  used  in  the  preparation  of  hydrogen 
a  few  drops  of  platinum  chloride  are  often  added  to  the 
solution  to  assist  the  chemical  action.  Nitric  acid  dissolves 
the  metal  readily,  with  the  formation  of  zinc  nitrate  and 
various  reduction  products  of  nitric  acid.  The  strong 
alkalis  act  upon  zinc  and  liberate  hydrogen : 

Zn  +  2  KOH  =  Zn(OK)2  +  2  H. 

The  product  of  this  reaction,  potassium  zincate,  is  a  salt  of 
zinc  hydroxide,  which  is  thus  seen  to  have  acid  properties, 
though  it  usually  acts  as  a  base. 

Uses  of  zinc.  The  metal  has  many  familiar  uses.  Rolled 
into  sheets,  it  is  used  as  a  lining  for  vessels  which  are  to 
contain  water.  As  a  thin  film  upon  the  surface  of  iron 
(galvanized  iron)  it  protects  the  iron  from  rust.  Iron  is 
usually  galvanized  by  dipping  it  into  a  bath  of  melted  zinc, 
but  electrical  methods  are  also  employed.  Zinc  plates  are 
used  in  many  forms  of  electrical  batteries.  In  the  labora- 
tory zinc  is  used  in  the  preparation  of  hydrogen,  and  in 
the  form  of  zinc  dust  as  a  reducing  agent. 

One  of  the  largest  uses  of  zinc  is  in  the  manufacture  of 
alloys.  Brass,  an  alloy  of  zinc  and  copper,  is  the  most 
important  of  these ;  German  silver,  consisting  of  copper, 
zinc,  and  nickel,  has  many  uses ;  various  bronzes,  coin 
metals,  and  bearing  metals  also  contain  zinc.  Its  ability 
to  alloy  with  silver  finds  application  in  the  separation  of 
silver  from  lead  (see  silver). 

Compounds  of  zinc.  In  general,  the  compounds  of  zinc  are 
similar  in  formula  and  appearance  to  those  of  magnesium, 


324     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

but  in  other  properties  they  often  differ  markedly.  A 
number  of  them  have  value  in  commercial  ways. 

Zinc  oxide  (zinc  white)  (ZnO).  Zinc  oxide  occurs  in 
impure  form  in  nature,  being  colored  red  by  manganese 
and  iron  compounds.  It  can  be  prepared  just  like  mag- 
nesium oxide,  but  is  more  often  made  by  burning  the 
metal. 

Zinc  oxide  is  a  pure  white  powder  which  becomes  yel- 
low on  heating  and  regains  its  white  color  when  cold.  It  is 
much  used  as  a  white  pigment  in  paints,  under  the  name 
of  zinc  white,  and  has  the  advantage  over  white  lead  in  that 
it  is  not  changed  in  color  by  sulphur  compounds,  while  lead 
turns  black.  It  is  also  used  in  the  manufacture  of  rubber 
goods. 

Commercial  preparation  of  zinc  oxide.  Commercially  it  is  often 
made  from  franklinite  in  the  following  way.  The  franklinite  is  mixed 
with  coal  and  heated  to  a  high  temperature  in  a  furnace,  by  which 
process  the  zinc  is  set  free  and  converted  into  vapor.  As  the  vapor 
leaves  the  furnace  through  a  conduit  it  meets  a  current  of  air  and 
takes  fire  in  it,  forming  zinc  oxide.  The  oxide  passes  on  and  is  fil- 
tered from  the  air  through  canvas  bags,  which  allow  the  air  to  pass 
but  retain  the  oxide.  It  is  thus  made  by  burning  the  metal,  though 
the  metal  is  not  actually  isolated  in  the  process. 

Soluble  salts.  The  soluble  salts  of  zinc  can  be  made  by 
dissolving  the  metal  or  the  oxide  in  the  appropriate  acid. 
They  are  all  somewhat  poisonous.  The  sulphate  and  chlo- 
ride are  the  most  familiar. 

Zinc  sulphate  (white  vitriol}  (ZnSO4-;  H2O).  This  salt 
is  readily  crystallized  from  strong  solutions  in  transparent 
colorless  crystals.  It  is  prepared  commercially  by  careful 
roasting  of  the  sulphide  : 

ZnS  +  4O-  ZnSO4. 


THE  MAGNESIUM  FAMILY  325 

Zinc  chloride  (ZnCl2  •  H2O).  When  a  solution  of  zinc 
chloride  is  slowly  evaporated  a  salt  of  the  composition 
ZnCl2  •  H2O  crystallizes  out.  If  the  water  is  completely 
expelled  by  heat  and  the  residue  distilled,  the  anhydrous 
chloride  is  obtained  and  may  be  cast  into  sticks  or  broken 
into  lumps.  In  this  distillation,  just  as  in  heating  mag- 
nesium chloride,  some  of  the  chloride  is  decomposed  : 

ZnCl2  •  H2O  =  ZnO  4-  2  HCL 

The  anhydrous  chloride  has  a  great  affinity  for  water,  and 
is  used  as  a  dehydrating  agent.  It  is  also  a  germicide,  and 
wood  which  is  to  be  exposed  to  conditions  which  favor  de- 
cay, as,  for  example,  railroad  ties,  is  often  soaked  in  solu- 
tions of  this  salt. 

Insoluble  compounds.  The  insoluble  compounds  of  zinc 
can  be  prepared  by  precipitation.  The  most  important  are 
the  sulphide,  carbonate,  and  hydroxide. 

Zinc  sulphide  (ZnS).  This  substance  occurs  as  the  min- 
eral sphalerite,  and  is  one  of  the  most  valued  ores  of  zinc. 
Very  large  deposits  occur  in  southwestern  Missouri.  The 
natural  mineral  is  found  in  large  crystals  or  masses,  resem- 
bling resin  in  color  and  luster.  When  prepared  by  precipi- 
tation the  sulphide  is  white. 

CADMIUM 

The  element.  This  element  occurs  in  small  quantities  in 
some  zinc  ores.  In  the  course  of  the  metallurgy  of  zinc 
the  cadmium  compounds  undergo  chemical  changes  quite 
similar  to  those  of  the  zinc  compounds,  and  the  cadmium 
distills  along  with  the  zinc.  Being  more  volatile,  it  comes 
over  with  the  first  of  the  zinc  and  is  prepared  from  the  first 
portions  of  the  distillate  by  special  methods  of  purification. 


326     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

The  element  very  closely  resembles  zinc  in  most  respects. 
Some  of  its  alloys  are  characterized  by  having  low  melting 
points. 

Compounds  of  cadmium.  Among  the  compounds  of  cad- 
mium may  be  mentioned  the  chloride  (CdCl2  •  2H2O),  the  sul- 
phate^ CdSO4-8  H2O),andthe  nitrate  (Cd(NO3)2 -4  H2O). 
These  are  white  solids  soluble  in  water.  The  sulphide 
(CdS)  is  a  bright  yellow  substance  which  is  insoluble  in 
water  and  in  dilute  acids.  It  is  valuable  as  a  pigment  in 
fine  paints. 

EXERCISES 

1.  What  properties  have  the  metals  of  the  magnesium  family  in 
common  with  the  alkali  metals ;  with  the  alkaline-earth  metals  ? 

2.  Compare  the  action  of  the  metals  of  the  magnesium  group  on 
water  with  that  of  the  other  metals  studied. 

3.  What  metals  already  studied  are  prepared  by  electrolysis  ? 

4.  Write  the  equations  representing  the  reactions  between  magne- 
sium and  hydrochloric  acid  ;  between  magnesium  and  dilute  sulphuric 
acid. 

5.  What  property  of  magnesium  was  taken  advantage  of  in  the 
isolation  of  argon  ? 

6.  With  phosphoric  acid  magnesium  forms  salts  similar  to  those 
of  calcium.    Write  the  names  and  formulas   of  the  corresponding 
magnesium  salts. 

7.  How  could  you  distinguish  between  magnesium  chloride  and 
magnesium  sulphate  ?  between  Glauber's  salts  and  Epsom  salts  ? 

8.  What  weight  of  carnallite  is  necessary  in  the  preparation  of 
500  g.  of  magnesium  ? 

9.  Account  for  the  fact  that  paints  made  of  zinc  oxide  are  not 
colored  by  hydrosulphuric  acid. 

10.  What  hydroxide  studied,  other  than  zinc  hydroxide,  has  both 
acid  and  basic  properties  ? 

11.  Write  equations  showing  how  the  following  compounds  of 
zinc  may  be  obtained  from  metallic  zinc  :  the  oxide,  chloride,  nitrate, 
carbonate,  sulphate,  sulphide,  hydroxide. 


CHAPTER    XXVI 
THE   ALUMINIUM  FAMILY 

The  family.  The  element  aluminium  is  the  most  abun- 
dant member  of  the  group  of  elements  known  as  the  alu- 
minium family  ;  indeed,  the  other  members  of  the  family  - 
gallium,  indium,  and  thallium  —  are  of  such  rare  occurrence 
that  they  need  not  be  separately  described.  The  elements 
of  the  family  are  ordinarily  trivalent,  so  that  the  formulas 
for  their  compounds  differ  from  those  of  the  elements  so 
far  studied.  Their  hydroxides  are  practically  insoluble  in 
water  and  are  very  weak  bases ;  indeed,  the  bases  are  so 
weak  that  their  salts  are  often  hydrolyzed  into  free  base  and 
free  acid  in  solution.  The  salts  formed  from  these  bases 
usually  contain  water  of  crystallization,  which  cannot  be 
driven  off  without  decomposing  them  more  or  less. 

The  trivalent  metals,  which  in  addition  to  aluminium 
include  also  iron  and  chromium,  are  sometimes  called  the 
earth  metals.  The  name  refers  to  the  earthy  appearance 
of  the  oxides  of  these  metals,  and  to  the  fact  that  many 
earths,  soils,  and  rocks  are  composed  in  part  of  these 
substances. 

ALUMINIUM 

Occurrence.  Aluminium  never  occurs  in  the  free  state 
in  nature,  owing  to  its  great  affinity  for  oxygen.  In  com- 
bined form,  as  oxides,  silicates,  and  a  few  other  salts,  it  is 
both  abundant  and  widely  distributed,  being  an  essential 

327 


328     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

constituent  of  all  soils  and  of  most  rocks  excepting  lime- 
stone and  sandstone.  Cryolite  (Na3AlF6),  found  in  Green- 
land, and  bauxite,  which  is  an  aluminium  hydroxide  usually 
mixed  with  some  iron  hydroxide,  are  important  minerals. 
It  is  estimated  that  aluminium  composes  about  8%  of  the 
earth's  crust.  In  the  industries  the  metal  is  called  alu- 
minum, but  its  chemical  name  is  aluminium. 


FIG. 


Preparation.  Aluminium  was  first  prepared  by  Wohler, 
in  1827,  by  heating  anhydrous  aluminium  chloride  with 
potassium:  A1C18  +  3  K  =  3  KC1  +  Al. 

This  method  was  tried  after  it  was  found  impossible  to 
reduce  the  oxide  of  aluminium  with  carbon.  The  metal 
possessed  such  interesting  properties  and  promised  to  be 
so  useful  that  many  efforts  were  made  to  devise  a  cheap 
way  of  preparing  it.  The  method  which  has  proved  most 
successful  consists  in  the  electrolysis  of  the  oxide  dissolved 
in  melted  cryolite. 

Metallurgy.  An  iron  box  A  (Fig.  82)  about  eight  feet  long  and 
six  feet  wide  is  connected  with  a  powerful  generator  in  such  a  way 


THE  ALUMINIUM  FAMILY  329 

as  to  serve  as  the  cathode  upon  which  the  aluminium  is  deposited. 
Three  or  four  rows  of  carbon  rods  B  dip  into  the  box  and  serve  as 
the  anodes.  The  box  is  partially  filled  with  cryolite  and  the  current 
is  turned  on,  generating  enough  heat  to  melt  the  cryolite.  Aluminium 
oxide  is  then  added,  and  under  the  influence  of  the  electric  current  it 
decomposes  into  aluminium  and  oxygen.  The  temperature  is  main- 
tained above  the  melting  point  of  aluminium,  and  the  liquid  metal, 
being  heavier  than  cryolite,  sinks  to  the  bottom  of  the  vessel,  from 
which  it  is  tapped  off  from  time  to  time  through  the  tap  hole  C.  The 
oxygen  in  part  escapes  as  gas,  and  in  part  combines  with  the  carbon 
of  the  anode,  the  combustion  being  very  brilliant.  The  process  is 
carried  on  at  Niagara  Falls. 

The  largest  expense  in  the  process,  apart  from  the  cost  of  elec- 
trical energy,  is  the  preparation  of  aluminium  oxide  free  from  other 
oxides,  for  most  of  the  oxide  found  in  nature  is  too  impure  to  serve 
without  refining.  Bauxite  is  the  principal  ore  used  as  a  source  of  the 
aluminium  because  it  is  converted  into  pure  oxide  without  great  diffi- 
culty. Since  common  clay  is  a  silicate  of  aluminium  and  is  every- 
where abundant,  it  might  be  expected  that  this  would  be  utilized  in 
the  preparation  of  aluminium.  It  is,  however,  very  difficult  to  extract 
the  aluminium  from  a  silicate,  and  no  practical  method  has  been 
found  which  will  accomplish  this. 

Physical  properties.  Aluminium  is  a  tin-white  metal 
which  melts  at  640°  and  is  very  light,  having  a  density  of 
2.68.  It  is  stiff  and  strong,  and  with  frequent  annealing 
can  be  rolled  into  thin  foil.  It  is  a  good  conductor  of  heat 
and  electricity,  though  not  so  good  as  copper  for  a  given 
cross  section  of  wire. 

Chemical  properties.  Aluminium  is  not  perceptibly  acted 
on  by  boiling  water,  and  moist  air  merely  dims  its  luster. 
Further  action  is  prevented  in  each  case  by  the  formation 
of  an  extremely  thin  film  of  oxide  upon  the  surface  of  the 
metal.  It  combines  directly  with  chlorine,  and  when  heated 
in  oxygen  burns  with  great  energy  and  the  liberation  of 
much  heat.  It  is  therefore  a  good  reducing  agent.  Hydro- 
chloric acid  acts  upon  it,  forming  aluminium  chloride ; 


330 


AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


nitric  acid  and  dilute  sulphuric  acid  have  almost  no  action 
on  it,  but  hot,  concentrated  sulphuric  acid  acts  upon  it  in 
the  same  way  as  upon  copper  : 

2  Al  +  6  H2SO4  =  A12(SO4)3  +  6  H2O  +  3  SO2. 

Alkalis  readily  attack  the  metal,  liberating  hydrogen,  as  in 
the  case  of  zinc  : 

Al  +  3  KOH  =  A1(OK)8  +  3  H. 

Salt  solutions,  such  as  sea  water,  corrode  the  metal  rapidly. 
It  alloys  readily  with  other  metals. 

Uses  of  aluminium.  These  properties  suggest  many  uses 
for  the  metal.  Its  lightness,  strength,  and  permanence  make 
it  well  adapted  for  many  construction  purposes.  These 
same  properties  have  led  to  its  extensive  use  in  the  manu- 
facture of  cooking  utensils.  The^fact  that  it  is  easily  cor- 
roded by  salt  solutions  is,  however,  a  disadvantage.  Owing 
to  its  small  resistance  to  electrical  currents,  it  is  replacing 
copper  to  some  extent  in  electrical  construction,  especially 
for  trolley  and  power  wires.  Some  of  its  alloys  have  very 
valuable  properties,  and  a  considerable  part  of  the  alumin- 
ium manufactured  is  used  for  this  purpose.  Aluminium 
bronze,  consisting  of  about  90%  copper  and  10%  aluminium, 
has  a  pure  golden  color,  is  strong  and  malleable,  is  easily 
cast,  and  is  permanent  in  the  air.  Considerable  amounts  of 
aluminium  steel  are  also  made. 

Goldschmidt  reduction  process.  Aluminium  is  frequently 
employed  as  a  powerful  reducing  agent,  many  metallic 
oxides  which  resist  reduction  by  carbon  being  readily 
reduced  by  it.  The  aluminium  in  the  form  of  a  fine 
powder  is  mixed  with  the  metallic  oxide,  together  with 
some  substance  such  as  fluorspar  to  act  as  a  flux.  The 
mixture  is  ignited,  and  the  aluminium  unites  with  the 


THE  ALUMINIUM  FAMILY  331 

oxygen  of  the  metallic  oxide,  liberating  the  metal.  This 
collects  in  a  fused  condition  under  the  flux. 

An  enormous  quantity  of  heat  is  liberated  in  this  reaction, 
and  a  temperature  as  high  as  3500°  can  be  reached.  The 
heat  of  the  reaction  is  turned  to  practical  account  in  weld- 
ing car  rails,  steel  castings,  and  in  similar  operations  where 
an  intense  local  heat  is  required.  A  mixture  of  aluminium 
with  various  metallic  oxides,  ready  prepared  for  such  pur- 
poses, is  sold  under  the  name  of  thermite. 

Preparation  of  chromium  by  the  Goldschmidt  method.  A  mixture  of 
chromium  oxide  and  aluminium  powder  is  placed  in  a  Hessian  cruci- 
ble (A,  Fig.  83),  and  on  top  of  it  is 
placed  a  small  heap  B  of  a  mixture 
of  sodium  peroxide  and  aluminium, 
into  which  is  stuck  a  piece  of  magne-  ~ 

sium  ribbon  C.  Powdered  fluorspar  D 
is  placed  around  the  sodium  peroxide, 
after  which  the  crucible  is  set  on  a 
pan  of  sand  and  the  magnesium  ribbon 
ignited.  When  the  flame  reaches  the 
sodium  peroxide  mixture  combustion  of 
the  aluminium  begins  with  almost  ex- 
plosive violence,  so  that  great  care  must  FIG  8_ 
be  taken  in  the  experiment.  The  heat 

of  this  combustion  starts  the  reaction  in  the  chromium  oxide  mixture, 
and  the  oxide  is  reduced  to  metallic  chromium.  When  the  crucible 
has  cooled  a  button  of  chromium  will  be  found  in  the  bottom. 

Aluminium  oxide  (A12O3).  This  substance  occurs  in 
several  forms  in  nature.  The  relatively  pure  crystals  are 
called  corundum,  while  emery  is  a  variety  colored  dark 
gray  or  black,  usually  with  iron  compounds.  In  transpar- 
ent crystals,  tinted  different  colors  by  traces  of  impurities, 
it  forms  such  precious  stones  as  the  sapphire,  oriental 
ruby,  topaz,  and  amethyst.  All  these  varieties  are  very 


332     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

hard,  falling  little  short  of  the  diamond  in  this  respect. 
Chemically  pure  aluminium  oxide  can  be  made  by  igniting 
the  hydroxide,  when  it  forms  an  amorphous  white  powder : 

2  A1(OH)3  =  A12O3  +  3  H2O. 

The  natural  varieties,  corundum  and  emery,  are  used  for 
cutting  and  grinding  purposes  ;  the  purest  forms,  together 
with  the  artificially  prepared  oxide,  are  largely  used  in  the 
preparation  of  aluminium. 

Aluminium  hydroxide  (A1(OH)8).  The  hydroxide  occurs 
in  nature  as  the  mineral  hydrargyllite,  and  in  a  partially 
dehydrated  form  called  bauxite.  It  can  be  prepared  by 
adding  ammonium  hydroxide  to  any  soluble  aluminium  salt, 
forming  a  semi-transparent  precipitate  which  is  insoluble  in 
water  but  very  hard  to  filter.  It  dissolves  in  most  acids  to 
form  soluble  salts,  and  in  the  strong  bases  to  form  alumi- 
nates,  as  indicated  in  the  equations 

A1(OH)3  +  3  HC1  =  A1C18  +  3  H2O, 
A1(OH)3  +  3  NaOH  =  Al(ONa)3  +  3  H2O. 

It  may  act,  therefore,  either  as  a  weak  base  or  as  a  weak 
acid,  its  action  depending  upon  the  character  of  the  sub- 
stances with  which  it  is  in  contact.  When  heated  gently 
the  hydroxide  loses  part  of  its  hydrogen  and  oxygen 
according  to  the  equation 

•  A1(OH)3  =  A1O  •  OH  +  H2O. 

This  substance,  the  formula  of  which  is  frequently  written 
HA1O2,  is  a  more  pronounced  acid  than  is  the  hydroxide, 
and  its  salts  are  frequently  formed  when  aluminium  com- 
pounds are  fused  with  alkalis.  The  magnesium  salt 
Mg(AlO2)2  is  called  spinel,  and  many  other  of  its  salts, 
called  aluminates,  are  found  in  nature. 


THE  ALUMINIUM  FAMILY  333 

When  heated  strongly  the  hydroxide  is  changed  into  oxide, 
which  will  not  again  take  up  water  on  being  moistened. 

Mordants  and  dyeing.  Aluminium  hydroxide  has  the  peculiar 
property  of  combining  with  many  soluble  coloring  materials  and 
forming  insoluble  products  with  them.  On  this  account  it  is  often 
used  as  a  filter  to  remove  objectionable  colors  from  water.  This 
property  also  leads  to  its  wide  use  in  the  dye  industry.  Many  dyes 
will  not  adhere  to  natural  fibers  such  as  cotton  and  wool,  that  is,  will 
not  "dye  fast."  If,  however,  the  cloth  to  be  dyed  is  soaked  in  a 
solution  of  aluminium  compounds  and  then  treated  with  ammonia, 
the  aluminium  salts  which  have  soaked  into  the  fiber  will  be  con- 
verted into  the  hydroxide,  which,  being  insoluble,  remains  in  the 
body  of  it.  If  the  fiber  is  now  dipped  into  a  solution  of  the  dye,  the 
aluminium  hydroxide  combines  with  the  color  material  and  fastens, 
or  "  fixes,"  it  upon  the  fiber.  A  substance  which  serves  this  purpose 
is  called  a  mordant,  and  aluminium  salts,  particularly  the  acetate, 
are  used  in  this  way. 

Aluminium  chloride  (A1C13  •  6  H2O).  This  substance 
is  prepared  by  dissolving  the  hydroxide  in  hydrochloric 
acid  and  evaporating  to  crystallization.  When  heated  it  is 
converted  into  the  oxide,  resembling  magnesium  in  this 
respect : 

2  (A1C18  •  6  H2O)  =  A12O8  +  6  HC1  +  9  H2O. 

The  anhydrous  chloride,  which  has  some  important  uses, 
is  made  by  heating  aluminium  turnings  in  a  current  of 
chlorine. 

Alums.  Aluminium  sulphate  can  be  prepared  by  the 
action  of  sulphuric  acid  upon  aluminium  hydroxide.  It  has 
the  property  of  combining  with  the  sulphates  of  the  alkali 
metals  to  form  compounds  called  alums.  Thus,  with  potas- 
sium sulphate  the  reaction  is  expressed  by  the  equation 

K2SO4  +  A12(SO4)3  +  24  H2O  -  2  (KA1(SO4)2-  12  H2O). 


334     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


Under  similar  conditions  ammonium  sulphate  yields  ammo- 
nium alum  : 

(NH4)2S04  +  A12(S04)3  +  24  H20  = 
2(NH4Al(SO4)2-i2H2O). 

Other  trivalent  sulphates  besides  aluminium  sulphate 
can  form  similar  compounds  with  the  alkali  sulphates, 
and  these  compounds  are  also  called  alums,  though  they 
contain  no  aluminium.  They  all  crystallize  in  octahedra 
and  contain  twelve  molecules  of  water  of  crystallization. 
The  alums  most  frequently  prepared  are  the  following : 


Potassium  alum  .  . 
Ammonium  alum  .  . 
Ammonium  iron  alum 
Potassium  chrome  alum 


KAl(SO4)2.i2H20. 
NH4A1(SO4)2.  i2H20. 
NH4Fe(SO4)2.  i2H20. 
KCr(SO4)8.i2H8Q. 


An  alum  may  therefore  be  regarded  as  a  compound  derived 
from  two  molecules  of  sulphuric  acid,  in 
which  one  hydrogen  atom  has  been  dis- 
placed by  the  univalent  alkali  atom,  and 
the  other  three  hydrogen  atoms  by  an 
atom  of  one  of  the  trivalent  metals,  such 
as  aluminium,  iron,  or  chromium. 


mm 


FIG. 84 


Very  large,  well-formed  crystals  of  an  alum 
can  be  prepared  by  suspending  a  small  crystal 
by  a  thread  in  a  saturated  solution  of  the  alum, 

as  shown  in  Fig.  84.    The  small  crystal  slowly  grows  and  assumes 

a  very  perfect  form. 

Other  salts  of  aluminium.  While  aluminium  hydroxide 
forms  fairly  stable  salts  with  strong  acids,  it  is  such  a  weak 
base  that  its  salts  with  weak  acids  are  readily  hydrolyzed. 
Thus,  when  an  aluminium  salt  and  a  soluble  carbonate  are 


THE  ALUMINIUM  FAMILY  335 

brought  together  in  solution  we  should  expect  to  have  alu- 
minium carbonate  precipitated  according  to  the  equation 

3  Na2C03  +  2  A1C18  -  A12(CO3)3  +  6  NaCL 

But  if  it  is  formed  at  all,  it  instantly  begins  to  hydrolyze, 
the  products  of  the  hydrolysis  being  aluminium  hydroxide 
and  carbonic  acid, 

A12(C03)3  +  6  H20  =  2  A1(OH)3  +  3  H2CO3. 

Similarly  a  soluble  sulphide,  instead  of  precipitating  alu- 
minium sulphide  (A12S8),  precipitates  aluminium  hydroxide  ; 
for  hydrogen  sulphide  is  such  a  weak  acid  that  the  alumin- 
ium sulphide  at  first  formed  hydrolyzes  at  once,  forming 
aluminium  hydroxide  and  hydrogen  sulphide  : 

3  Na2S  +  2.  A1C18  +  6  H2O  =  2.  A1(OH)3  +  6  NaCl  +  3  H2S. 

Alum  baking  powders.  It  is  because  of  the  hydrolysis  of  aluminium 
carbonate  that  alum  is  used  as  a  constituent  of  some  baking  powders. 
The  alum  baking  powders  consist  of  a  mixture  of  alum  and  sodium 
hydrogen  carbonate.  When  water  is  added  the  two  compounds 
react  together,  forming  aluminium  carbonate,  which  hydrolyzes  into 
aluminium  hydroxide  and  carbonic  acid.  The  carbon  dioxide  from 
the  latter  escapes  through  the  dough  and  in  so  doing  raises  it  into 
a  porous  condition,  which  is  the  end  sought  in  the  use  of  a  baking 
powder. 

Aluminium  silicates.  One  of  the  most  common  constitu- 
ents of  rocks  is  feldspar  (KAlSi3O8),  a  mixed  salt  of  potas- 
sium and  aluminium  with  the  polysilicic  acid  (H4Si3O8). 
Under  the  influence  of  moisture,  carbon  dioxide,  and 
changes  of  temperature  this  substance  is  constantly  being 
broken  down  into  soluble  potassium  compounds  and  hy- 
drated  aluminium  silicate.  This  compound  has  the  formula 
Al2Si2O7  •  2  H2O.  In  relatively  pure  condition  it  is  called 
kaolin;  in  the  impure  state,  mixed  with  sand  and  other 


336     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

substances,  it  forms  common  clay.  Mica  is  another  very 
abundant  mineral,  having  varying  composition,  but  being 
essentially  of  the  formula  KAlSiO4.  Serpentine,  talc, 
asbestos,  and  meerschaum  are  important  complex  silicates 
of  aluminium  and  magnesium,  and  granite  is  a  mechanical 
mixture  of  quartz,  feldspar,  and  mica. 

Ceramic  industries.  Many  articles  of  greatest  practical  importance, 
ranging  from  the  roughest  brick  and  tile  to  the  finest  porcelain  and 
chinaware,  are  made  from  some  form  of  kaolin,  or  clay.  No  very  pre- 
cise classification  of  such  ware  can  be  made,  as  the  products  vary 
greatly  in  properties,  depending  upon  the  materials  used  and  the  treat- 
ment during  manufacture. 

Porcelain  is  made  from  the  purest  kaolin,  to  which  must  be  added 
some  less  pure,  plastic  kaolin,  since  the  pure  substance  is  not  suffi- 
ciently plastic.  There  is  also  added  some  more  fusible  substance,  such 
as  feldspar,  gypsum,  or  lime,  together  with  some  pure  quartz.  The  con- 
stituents must  be  ground  very  fine,  and  when  thoroughly  mixed  and 
moistened  must  make  a  plastic  mass  which  can  be  molded  into  any 
desired  form.  The  article  molded  from  such  materials  is  then 
burned.  In  this  process  the  article  is  slowly  heated  to  a  point  at 
which  it  begins  to  soften  and  almost  fuse,  and  then  it  is  allowed  to 
cool  slowly.  At  this  stage,  a  very  thin  vessel  will  be  translucent 
and  have  an  almost  glassy  fracture  ;  if,  however,  it  is  somewhat  thicker, 
or  has  not  been  heated  quite  so  high,  it  will  still  be  porous,  and 
partly  on  this  account  and  partly  to  improve  its  appearance  it  is 
usually  glazed. 

Glazing  is  accomplished  by  spreading  upon  the  object  a  thin  layer 
of  a  more  fusible  mixture  of  the  same  materials  as  compose  the  body 
of  the  object  itself,  and  again  heating  until  the  glaze  melts  to  a  trans- 
parent glassy  coating  upon  the  surface  of  the  vessel.  In  some  cases 
fusible  mixtures  of  quite  different  composition  from  that  used  in  fash- 
ioning the  vessel  may  be  used  as  a  glaze.  Oxides  of  lead,  zinc,  and 
barium  are  often  used  in  this  way. 

When  less  carefully  selected  materials  are  used,  or  quite  thick 
vessels  are  made,  various  grades  of  stoneware  are  produced.  The 
inferior  grades  are  glazed  by  throwing  a  quantity  of  common  salt 
into  the  kiln  towards  the  end  of  the  first  firing.  In  the  form  of  vapor 


THE  ALUMINIUM  FAMILY  337 

the  salt  attacks  the  surface  of  the  baked  ware  and  forms  an  easily 
fusible  sodium  silicate  upon  it,  which  constitutes  a  glaze. 

Vitrified  bricks,  made  from  clay  or  ground  shale,  are  burned  until 
the  materials  begin  to  fuse  superficially,  forming  their  own  glaze. 
Other  forms  of  brick  and  tile  are  not  glazed  at  all,  but  are  left  porous. 
The  red  color  of  ordinary  brick  and  earthenware  is  due  to  an  oxide 
of  iron  formed  in  the  burning  process. 

The  decorations  upon  china  are  sometimes  painted  upon  the  baked 
ware  and  then  glazed  over,  and  sometimes  painted  upon  the  glaze  and 
burned  in  by  a  third  firing.  Care  must  be  taken  to  use  such  pigments 
as  are  not  affected  by  a  high  heat  and  do  not  react  chemically  with 
the  constituents  of  the  baked  ware  or  the  glaze. 


EXERCISES 

1.  What  metals  and  compounds  studied  are  prepared  by  electrol- 
ysis? 

2.  Write  the  equation  for  the  reaction  between  aluminium  and 
hydrochloric  acid  ;  between  aluminium  and  sulphuric  acid  (in  two 
steps). 

3.  What  hydroxides  other  than  aluminium  hydroxide  have  both 
acid  and  basic  properties  ? 

4.  Write    equations   showing   the   methods   used    for   preparing 
aluminium  hydroxide  and  sulphate. 

5.  Write  the  general  formula  of  an  alum,  representing  an  atom  of 
an  alkali  metal  by  X  and  an  atom  of  a  trivalent  metal  by  Y. 

6.  What  is  meant  by  the  term  poly  silicic  acid,  as  used  in  the  dis- 
cussion of  aluminium  silicates  ? 

7.  Compare  the  properties  of   the  hydroxides  of   the  different 
groups  of  metals  so  far  studied. 

8.  In  what  respects  does  aluminium  oxide   differ  from  calcium 
oxide  in  properties  ? 

9.  Supposing  bauxite  to  be  90%  aluminium  hydroxide,  what  weight 
of  it  is  necessary  for  the  preparation  of  i  oo  kg.  of  aluminium  ? 


CHAPTER    XXVII 
THE  IRON  FAMILY 


APPROXIMATE 

SYMBOL 

ATOMIC 
WEIGHT 

DENSITY 

MELTING 

OXIDES 

POINT 

Iron 

Fe 

CC.Q 

7.Q-? 

1800° 

FeO,   Fe2O3 

Cobalt     .... 

Co 

59-o 

8.55 

1800° 

CoO,  Co2O3 

Nickel     .... 

Ni 

58.7 

8.9 

1600° 

NiO,  Ni2O3 

The  family.  The  elements  iron,  cobalt,  and  nickel  form 
a  group  in  the  eighth  column  of  the  periodic  table.  The 
atomic  weights  of  the  three  are  very  close  together,  and 
there  is  not  the  same  gradual  gradation  in  the  properties 
of  the  three  elements  that  is  noticed  in  the  families  in  which 
the  atomic  weights  differ  considerably  in  magnitude.  The 
elements  are  very  similar  in  properties,  the  similarity  being 
so  great  in  the  case  of  nickel  and  cobalt  that  it  is  difficult 
to  separate  them  by  chemical  analysis. 

The  elements  occur  in  nature  chiefly  as  oxides  and  sul- 
phides, though  they  have  been  found  in  very  small  quan- 
tities in  the  native  state,  usually  in  meteorites.  Their 
sulphides,  carbonates,  and  phosphates  are  insoluble  in  water, 
the  other  common  salts  being  soluble.  Their  salts  are  usu- 
ally highly  colored,  those  of  iron  being  yellow  or  light  green 
as  a  rule,  those  of  nickel  darker  green,  while  cobalt  salts  are 
usually  rose  colored.  The  metals  are  obtained  by  reducing 
the  oxides  with  carbon. 

338 


THE  IRON  FAMILY  339 

IRON 

Occurrence.  The  element  iron  has  long  been  known,  since 
its  ores  are  very  abundant  and  it  is  not  difficult  to  prepare 
the  metal  from  them  in  fairly  pure  condition.  It  occurs  in 
nature  in  many  forms  of  combination,  —  in  large  deposits 
as  oxides,  sulphides,  and  carbonates,  and  in  smaller  quan- 
tities in  a  great  variety  of  minerals.  Indeed,  very  few 
rocks  or  soils  are  free  from  small  amounts  of  iron,  and  it  is 
assimilated  by  plants  and  animals  playing  an  important 
part  in  life  processes. 

Metallurgy.  It  will  be  convenient  to  treat  of  the  metal- 
lurgy of  iron  under  two  heads,  —  Materials  Used  and 
Process. 

Materials  used.  Four  distinct  materials  are  used  in  the 
metallurgy  of  iron  : 

1 .  Iron  ore.    The  ores  most  frequently  used  in  the  metal- 
lurgy of  iron  are  the  following : 

Hematite Fe2O3. 

Magnetite Fe3O4. 

Siderite  FeCO3. 

Limonite .  2  Fe2O3 . 3  H2O. 

These  ores  always  contain  impurities,  such  as  silica,  sul- 
phides, and  earthy  materials.  All  ores,  with  the  exception 
of  the  oxides,  are  first  roasted  to  expel  any  water  and  carbon 
dioxide  present  and  to  convert  any  sulphide  into  oxide. 

2.  Carbon.    Carbon  in  some  form  is  necessary  both  as 
a  fuel  and  as  a  reducing  agent.     In  former  times  wood 
charcoal  was  used  to  supply  the  carbon,  but  now  anthracite 
coal  or  coke  is  almost  universally  used. 

3.  Hot  air.    To  maintain  the  high  temperature  required 
for  the  reduction  of  iron  a  very  active  combustion  of  fuel 


340     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

is  necessary.  This  is  secured  by  forcing  a  strong  blast  of 
hot  air  into  the  lower  part  of  the  furnace  during  the  re- 
duction process. 

4.  Flux,  (a)  Purpose  of  the  flux.  All  the  materials  which 
enter  the  furnace  must  leave  it  again  either  in  the  form  of 
gases  or  as  liquids.  The  iron  is  drawn  off  as  the  liquid 
metal  after  its  reduction.  To  secure  the  removal  of  the 
earthy  matter  charged  into  the  furnace  along  with  the  ore, 
materials  are  added  to  the  charge  which  will,  at  the  high 
temperature  of  the  furnace,  combine  with  the  impurities  in 
the  ore,  forming  a  liquid.  The  material  added  for  this  pur- 
pose is  called  the  flux ;  the  liquid  produced  from  the  flux 
and  the  ore  is  called  slag. 

(b)  Function  of  the  slag.    While   the   main  purpose  of 
adding  flux  to  the  charge  is  to  remove  from  the  furnace  in 
the  form  of  liquid  slag  the  impurities  originally  present  in 
the  ore,  the  slag  thus  produced  serves  several  other  func- 
tions.   It  keeps  the  contents  of  the  furnace  in  a  state  of 
fusion,  thus  preventing  clogging,  and  makes  it  possible  for 
the  small  globules  of  iron  to  run  together  with  greater  ease 
into  one  large  liquid  mass. 

(c)  Character  of  the  slag.    The  slag  is  really  a  kind  of 
readily  fusible  glass,  being  essentially  a  calcium-aluminium 
silicate.    The  ore  usually  contains  silica  and  some  alumin- 
ium compounds,  so  that  limestone   (which    also   contains 
some  silica  and  aluminium)  is  added  to  furnish  the  calcium 
required  for  the  slag.    If  the  ore  and  the  limestone  do  not 
contain  a  sufficient  amount  of  silica  and  aluminium  for  the 
formation  of  the  slag,  these  ingredients  are  added  in  the 
form  of  sand  and  feldspar.    In  the  formation  of  slag  from 
these  materials  the  ore  is  freed  from  the  silica  and  alu- 
minium which  it  contained. 


THE  IRON  FAMILY 


341 


Process.  The  reduction  of  iron  is  carried  out  in  large 
towers  called  blast  furnaces.  The  blast  furnace  (Fig.  85) 
is  usually  about  80  ft.  high  and  20  ft.  in  internal  diameter 
at  its  widest  part,  narrowing  somewhat  both  toward  the 
top  and  toward  the  bottom.  The  walls  are  built  of  steel 
and  lined  with  fire  brick.  The  base  is  provided  with 
a  number  of  pipes  7",  called  tuyers, 
through  which  hot  air  can  be  forced 
into  the  furnace.  The  tuyers  are  sup- 
plied from  a  large  pipe  S,  which  circles 
the  furnace  as  a  girdle.  The  base  has 
also  an  opening  M,  through  which  the 
liquid  metal  can  be  drawn  off  from 
time  to  time,  and  a  second  opening  P,  D 
somewhat  above  the  first,  through 
which  the  excess  of  slag  overflows. 
The  top  is  closed  by  a  movable  trap 
£7 and  C'y  called  the  cone,  and  through 
this  the  materials  to  be  used  are  in- 
troduced. The  gases  produced  by  s 
the  combustion  of  the  fuel  and  the  re- 
duction of  the  ore,  together  with  the 
nitrogen  of  the  air  forced  in  through 
the  tuyers,  escape  through  pipes  D, 


FIG. 85 


called  downcomer  pipes,  which  leave  the  furnace  near  the 
top.  These  gases  are  very  hot  and  contain  combustible 
substances,  principally  carbon  monoxide ;  they  are  there- 
fore utilized  as  fuel  for  the  engines  and  also  to  heat  the  blast 
admitted  through  the  tuyers.  The  lower  part  of  the  fur- 
nace is  pften  furnished  with  a  water  jacket.  This  consists 
of  a  series  of  pipes  W  built  into  the  walls,  through  which 
water  can  be  circulated  to  reduce  their  temperature. 


342     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Charges  consisting  of  coke  (or  anthracite  coal),  ore,  and 
flux  in  proper  proportions  are  introduced  into  the  furnace 
at  intervals  through  the  trap  top.  The  coke  burns  fiercely 
in  the  hot-air  blast,  giving  an  intense  heat  and  forming 
carbon  monoxide.  The  ore,  working  down  in  the  furnace 
as  the  coke  burns,  becomes  very  hot,  and  by  the  combined 
reducing  action  of  the  carbon  and  carbon  monoxide  is 
finally  reduced  to  metal  and  collects  as  a  liquid  in  the  bot- 
tom of  the  furnace,  the  slag  floating  on  the  molten  iron. 
After  a  considerable  amount  of  the  iron  has  collected  the 
slag  is  drawn  off  through  the  opening  P.  The  molten  iron 
is  then  drawn  off  into  large  ladles  and  taken  to  the  con- 
verters for  the  manufacture  of  steel,  or  it  is  run  out  into 
sand  molds,  forming  the  bars  or  ingots  called  "  pigs."  The 
process  is  a  continuous  one,  and  when  once  started  it  is  kept 
in  operation  for  months  or  even  years  without  interruption. 

It  seems  probable  that  the  first  product  of  combustion  of  the 
carbon,  at  the  point  where  the  tuyers  enter  the  furnace,  is  carbon 
dioxide.  This  is  at  once  reduced  to  carbon  monoxide  by  the  intensely 
heated  carbon  present,  so  that  no  carbon  dioxide  can  be  found  at 
that  point.  For  practical  purposes,  therefore,  we  may  consider  that 
carbon  monoxide  is  the  first  product  of  combustion. 

Varieties  of  iron.  The  iron  of  commerce  is  never  pure, 
but  contains  varying  amounts  of  other  elements,  such  as 
carbon,  silicon,  phosphorus,  sulphur,  and  manganese.  These 
elements  may  either  be  alloyed  with  the  iron  or  may  be 
combined  with  it  in  the  form  of  definite  chemical  com- 
pounds. In  some  instances,  as  in  the  case  of  graphite,  the 
mixture  may  be  merely  mechanical. 

The  properties  of  iron  are  very  much  modified  by  the 
presence  of  these  elements  and  by  the  form  of  the  combi- 
nation between  them  and  the  iron ;  the  way  in  which  the 


THE  IRON  FAMILY  343 

metal  is  treated  during  its  preparation  has  also  a  narked 
influence  on  its  properties.  Owing  to  these  facts  many 
kinds  of  iron  are  recognized  in  commerce,  the  chief  varie- 
ties  being  cast  iron,  wrought  iron,  and  steel. 

Cast  iron.  The  product  of  the  blast  furnace,  prepared  as 
just  described,  is  called  cast  iron.  It  varies  considerably 
in  composition,  usually  containing  from  90  to  95%  iron, 
the  remainder  being  largely  carbon  and  silicon  with  smaller 
amounts  of  phosphorus  and  sulphur.  When  the  melted 
metal  from  the  blast  furnace  is  allowed  to  cool  rapidly 
most  of  the  carbon  remains  in  chemical  combination  with 
the  iron,  and  the  product  is  called  white  cast  iron.  If  the 
cooling  goes  on  slowly,  the  carbon  partially  separates  as 
flakes  of  graphite  which  remain  scattered  through  the 
metal.  This  product  is  softer  and  darker  in  color  and  is 
called  gray  cast  iron. 

Properties  of  cast  iron.  Cast  iron  is  hard,  brittle,  and 
rather  easily  melted  (melting  point  about  1 100°).  It  can- 
not be  welded  or  forged  into  shape,  but  is  easily  cast  in 
sand  molds.  It  is  strong  and  rigid  but  not  elastic.  It  is 
used  for  making  castings  and  in  the  manufacture  of  other 
kinds  of  iron.  Cast  iron,  which  contains  the  metal  manga- 
nese up  to  the  extent  of  20%,  together  with  about  3% 
carbon,  is  called  spiegel  iron ;  when  more  than  this 
amount  of  manganese  is  present  the  product  is  called  fer- 
romanganese.  The  ferromanganese  may  contain  as  much 
as  80%  manganese.  These  varieties  of  cast  iron  are  much 
used  in  the  manufacture  of  steel. 

Wrought  iron.  Wrought  iron  is  made  by  burning  out 
from  cast  iron  most  of  the  carbon,  silicon,  phosphorus,  and 
sulphur  which  it  contains.  The  process  is  called  puddling, 
and  is  carried  out  in  a  furnace  constructed  as  represented 


344 


AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


in  Fig.  86.  The  floor  of  the  furnace  F  is  somewhat  con- 
cave and  is  made  of  iron  covered  with  a  layer  of  iron  oxide. 
A  long  flame  produced  by  burning  fuel  upon  the  grate  G  is 
directed  downward  upon  the  materials  placed  upon  the 
floor,  and  the  draught  is  maintained  by  the  stack  5.  A  is 
the  ash  box  and  T  a  trap  to  catch  the  solid  particles  car- 
ried into  the  stack  by  the  draught.  Upon  the 
flooE-af~the  furnace  is  placed  the  charge  of  cast 
iron,  together  with  a  small  amount  of  material 

to  make  a  slag.  The 
iron  is  soon  melted  by 
the  flame  directed  up- 
on it,  and  the  sulphur. 
phosphorus,  and  silicon 
are  oxidized  by  the  iron 
oxide,  forming  oxides 
which  are  anhydrides  of 
acids.  These  combine 
with  the  flux,  which 
is  basic  in  character,  or  with  the  iron  oxide,  to  form  a  slag. 
The  carbon-is  also  oxidized  and  escapes  as  carbon  dioxide. 
As  the  iron  is  freed  from  other  elements  it  becomes  pasty, 
owing  to  the  higher  melting  point  of  the  purer  iron,  and  in 
this  condition  forms  small  lumps  which  are  raked  together 
into  a  larger  one.  The  large  lump  is  then  removed  from 
the  furnace  and  rolled  or  hammered  into  bars,  the  slag 
being  squeezed  out  in  this  process.  The  product  has  a 
stranded  or  fibrous  structure.  The  product  of  a  puddling 
furnace  is  called  wrought  iron. 

Properties  of  wrought  iron.  Wrought  iron  is  nearly 
pure  iron,  usually  containing  about  o/j$£  of  other  sub- 
stances, chiefly  carbon.  It  is  tough,  malleable,  and  fibrous 

^- 


FIG.  86 


THE  IRON  FAMILY 


345 


in  structure.  It  is  easily  bent  and  is  not  elastic,  so  i^will  not 
sustain  pressure  as  well  as  cast  iron.  It  can  be  drawn  out 
into  wire  of  great  tensile  strength,  and  can  also  be  rolled 
into  thin  sheets  (sheet  iron).  It  melts  at  a  high  temper- 
ature (about  1600°)  and  is  therefore  forged  into  shape 
rather  than  cast.  If  melted,  it  would  lose  its  fibrous  struc- 
ture and  be  changed  into  a  low  carbon  steel. 

Steel.  Steel,  like  wrought  iron,  is  made  by  burning  out 
from  cast  iron  a  part  of  the  carbon,  silicon,  phosphorus,  and 
sulphur  which  it  contains ;  but  the  process  is  carried  out 
in  a  very  different  way,  and  usually,  though  not  always, 
more  carbon  is  found  in  steel  than  in  wrbught^jron..  A 
number  of  processes 
are  in  use,  but  nearly 
all  the  steel  of  com- 
merce is  made  by 
one  of  the  two  fol- 
lowing methods. 

I.  Bessemer  pro- 
cess. This  process, 
invented  about  1860, 
is  by  far  the  most 
important.  It  is  car- 
ried out  in  great  egg- 
shaped  crucibles 
called  converters  (Fig.  87),  each  one  of  which  will  hold  as 
much  as  15  tons  of  steel.  The  converter  is  built  of  steel 
and  lined  with  silica.  It  is  mounted  on  trunnions  T,  so 
that  it  can  be  tipped  over  on  its  side  for  filling  and  empty- 
ing. One  of  the  trunnions  is  hollow  and  a  pipe  P  connects 
it  with  an  air  chamber  .A,  which  forms  a  false  bottom  to 
the  converter.  The  true  bottom  is  perforated,  so  that  air 


FIG 


346     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

can  be  forced  in  by  an  air  blast  admitted  through  the 
trunnion  and  the  air  chamber. 

White-hot,  liquid  cast  iron  from  a  blast  furnace  is  run 
into  the  converter  through  its  open  necklike  top  O,  the 
converter  being  tipped  over  to  receive  it ;  the  air  blast  is 
then  turned  on  and  the  converter  rotated  to  a  nearly  vertical 
position.  The  elements  in  the  iron  are  rapidly  oxidized, 
the  silicon  first  and  then  the  carbon.  The  heat  liberated 
in  the  oxidation,  largely  due  to  the  combustion  of  silicon, 
keeps  the  iron  in  a  molten  condition.  When  the  carbon  is 
practically  all  burned  out  cast  iron  or  spiegel  iron,  con- 
taining a  known  percentage  of  carbon,  is  added  and  allowed 
to  mix  thoroughly  with  the  fluid.  The  steel  is  then  run 
into  molds,  and  the  ingots  so  formed  are  hammered  or 
rolled  into  rails  or  other  forms.  By  this  process  any 
desired  percentage  of  carbon  can  be  added  to  the  steel. 
Low  carbon  steel,  which  does  not  differ  much  from  wrought 
iron  in  composition,  is  now  made  in  this  way  and  is  replac- 
ing the  more  expensive  wrought  iron  for  many  purposes. 

The  basic  lining  process.  When  the  cast  iron  contains  phosphorus 
and  sulphur  in  appreciable  quantities,  the  lining  of  the  converter  is 
made  of  dolomite.  The  silicon  and  carbon  burn,  followed  by  the 
phosphorus  and  sulphur,  and  the  anhydrides  of  acids  so  formed  com- 
bine with  the  basic  oxides  of  the  lining,  forming  a  slag.  This  is 
known  as  the  basic  lining  process. 

2.  Open-hearth  process.  In  this  process  a  furnace  very 
similar  to  a  puddling  furnace  is  used,  but  it  is  lined  with 
silica  or  dolomite  instead  of  iron  oxide.  A  charge  consist- 
ing in  part  of  old  scrap  iron  of  any  kind  and  in  part  of  cast 
iron  is  melted  in  the  furnace  by  a  gas  flame.  The  silicon 
and  carbon  are  slowly  burned  away,  and  when  a  test  shows 
that  the  desired  percentage  of  carbon  is  present  the  steel 


THE  IRON  FAMILY  347 

is  run  out  of  the  furnace.  Steel  may  therefore  be  defined 
as  the  product  of  the  Bessemer  or  open-hearth  processes. 

Properties  of  steel.  Bessemer  and  open-hearth  steel  usu- 
ally contain  only  a  few  tenths  of  a  per  cent  of  carbon, 
less  than  o.i  %  silicon,  and  a  very  much  smaller  quantity  of 
phosphorus  and  sulphur.  Any  considerable  amount  of  the 
latter  elements  makes  the  steel  brittle,  the  sulphur  affect- 
ing it  when  hot,  and  the  phosphorus  when  cold.  This 
kind  of  steel  is  used  for  structural  purposes,  for  rails,  and 
for  nearly  all  large  steel  articles.  It  is  hard,  malleable, 
ductile,  and  melts  at  a  lower  temperature  than  wrought 
iron.  It  can  be  forged  into  shape,  rolled  into  sheets,  or 
cast  in  molds. 

Relation  of  the  three  varieties  of  iron.  It  will  be  seen  that 
wrought  iron  is  usually  very  nearly  pure  iron,  while  steel 
contains  an  appreciable  amount  of  alloy  material,  chiefly 
carbon,  and  cast  iron  still  more  of  the  same  substances.  It 
is  impossible,  however,  to  assign  a  given  sample  of  iron  to 
one  of  these  three  classes  on  the  basis  of  its  chemical 
composition  alone.  A  low  carbon  steel,  for  example,  may 
contain  less  carbon  than  a  given  sample  of  wrought  iron. 
The  real  distinction  between  the  three  is  the  process  by 
which  they  are  made.  The  product  of  the  blast  furnace  is 
cast  iron ;  that  of  the  puddling  furnace  is  wrought  iron ; 
that  of  the  Bessemer  and  open-hearth  methods  is  steel. 

Tool  steel.  Steel  designed  for  use  in  the  manufacture  of 
edged  tools  and  similar  articles  should  be  relatively  free 
from  silicon  and  phosphorus,  but  should  contain  from  0.5 
to  1.5%  carbon.  The  percentage  of  carbon  should  be  regu- 
lated by  the  exact  use  to  which  the  steel  is  to  be  put.  Steel 
of  this  character  is  usually  made  in  small  lots  from  either 
Bessemer  or  open-hearth  steel  in  the  following  way. 


348     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

A  charge  of  melted  steel  is  placed  in  a  large  crucible  and 
the  calculated  quantity  of  pure  carbon  is  added.  The  car- 
bon dissolves  in  the  steel,  and  when  the  solution  is  com- 
plete the  metal  is  poured  out  of  the  crucible.  This  is 
sometimes  called  crucible  steel. 

Tempering  of  steel.  Steel  containing  from  0.5  to  1.5% 
carbon  is  characterized  by  the  property  of  "taking  temper." 
When  the  hot  steel  is  suddenly  cooled  by  plunging  it  into 
water  or  oil  it  becomes  very  hard  and  brittle.  On  carefully 
reheating  this  hard  form  it  gradually  becomes  less  brittle 
and  softer,  so  that  by  regulating  the  temperature  to  which 
steel  is  reheated  in  tempering  almost  any  condition  of 
temper  demanded  for  a  given  purpose,  such  as  for  making 
springs  or  cutting  tools,  can  be  obtained. 

Steel  alloys.  It  has  been  found  that  small  quantities  of 
a  number  of  different  elements  when  alloyed  with  steel 
very  much  improve  its  quality  for  certain  purposes,  each 
element  having  a  somewhat  different  effect.  Among  the 
elements  most  used  in  this  connection  are  manganese,  sili- 
con, chromium,  nickel,  tungsten,  and  molybdenum. 

The  usual  method  for  adding  these  elements  to  the  steel 
is  to  first  prepare  a  very  rich  alloy  of  iron  with  the  element 
to  be  added,  and  then  add  enough  of  this  alloy  to  a  large 
quantity  of  the  steel  to  bring  it  to  the  desired  composition. 
A  rich  alloy  of  iron  with  manganese  or  silicon  can  be  pre- 
pared directly  in  a  blast  furnace,  and  is  called  ferro- 
manganese  or  ferrosilicon.  Similar  alloys  of  iron  with  the 
other  elements  mentioned  are  made  in  an  electric  furnace 
by  reducing  the  mixed  oxides  with  carbon. 

Pure  iron.  Perfectly  pure  iron  is  rarely  prepared  and  is 
not  adapted  to  commercial  uses.  It  can  be  made  by 
reducing  pure  oxide  of  iron  in  a  current  of  hydrogen  at  a 


THE  IRON  FAMILY 


349 


high  temperature.  Prepared  in  this  way  it  forms  a  black 
powder ;  when  melted  it  forms  a  tin-white  metal  which  is 
less  fusible  and  more  malleable  than  wrought  iron.  It  is 
easily  acted  upon  by  moist  air. 

Compounds  of  iron.  Iron  differs  from  the  metals  so  far 
studied  in  that  it  is  able  to  form  two  series  of  compounds 
in  which  the  iron  has  two  different  valences.  In  the  one 
series  the  iron  is  divalent  and  forms  compounds  which  in 
formulas  and  many  chemical  properties  are  similar  to  the 
corresponding  zinc  compounds.  It  can  also  act  as  a  triva- 
lent  metal,  and  in  this  condition  forms  salts  similar  to  those 
of  aluminium.  Those  compounds  in  which  the  iron  is  diva- 
lent are  known  as  ferrous  compounds,  while  those  in  which 
it  is  trivalent  are  known  as  ferric. 

Oxides  of  iron.  Iron  forms  several  oxides.  Ferrous  oxide 
(FeO)  is  not  found  in  nature,  but  can  be  prepared  artificially 
in  the  form  of  a  black  powder  which  easily  takes  up  oxygen, 
forming  ferric  oxide  : 

2  FeO  +  O  =  Fe2O3. 

Ferric  oxide  is  the  most  abundant  ore  of  iron  and  occurs 
in  great  deposits,  especially  in  the  Lake  Superior  region. 
It  is  found  in  many  mineral  varieties  which  vary  in  density 
and  color,  the  most  abundant  being  hematite,  which  ranges 
in  color  from  red  to  nearly  black.  When  prepared  by 
chemical  processes  it  forms  a  red  powder  which  is  used  as 
a  paint  pigment  (Venetian  red)  and  as  a  polishing  powder 
(rouge). 

Magnetite  has  the  formula  Fe3O4  and  is  a  combination 
of  FeO  and  Fe2O3.  It  is  a  very  valuable  ore,  but  is  less 
abundant  than  hematite.  It  is  sometimes  called  magnetic 
oxide  of  iron,  or  lodestone,  since  it  is  a  natural  magnet. 


350     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Ferrous  salts.  These  salts  are  obtained  by  dissolving 
iron  in  the  appropriate  acid,  or,  when  insoluble,  by  precipi- 
tation. They  are  usually  light  green  in  color  and  crystallize 
well.  In  chemical  reactions  they  are  quite  similar  to  the 
salts  of  magnesium  and  zinc,  but  differ  from  them  in  one 
important  respect,  namely,  that  they  are  easily  changed 
into  compounds  in  which  the  metal  is  trivalent.  Thus 
ferrous  chloride  treated  with  chlorine  or  aqua  regia  is 
changed  into  ferric  chloride  : 

FeCl2  +  Cl  =  FeCl3. 

Ferrous  hydroxide  exposed  to  moist  air  is  rapidly  changed 
into  ferric  hydroxide  : 

2  Fe(OH)2  +  H2O  +  O  =  2  Fe(OH)3. 

Ferrous  sulphate  (copperas,  green  vitriol)  (FeSO4  •  7  H2O). 
Ferrous  sulphate  is  the  most  familiar  ferrous  compound. 
It  is  prepared  commercially  as  a  by-product  in  the  steel- 
plate  mills.  Steel  plates  are  cleaned  by  the  action  of  dilute 
sulphuric  acid  upon  them,  and  in  the  process  some  of  the 
iron  dissolves.  The  liquors  are  concentrated  and  the  green 
vitriol  separates  from  them. 

Ferrous  sulphide  (FeS).  Ferrous  sulphide  is  sometimes 
found  in  nature  as  a  golden-yellow  crystalline  mineral.  It 
is  formed  as  a  black  precipitate  when  a  soluble  sulphide 
and  an  iron  salt  are  brought  together  in  solution : 

FeSO4  +  Na2S  =  FeS  +  Na2SO4. 

It  can  also  be  made  as  a  heavy  dark -brown  solid  by  fusing 
together  the  requisite  quantities  of  sulphur  and  iron.  It  is 
obtained  as  a  by-product  in  the  metallurgy  of  lead  : 

PbS  +  Fe  =  FeS  +  Pb. 


THE  IRON  FAMILY  351 

It  is  used  in  the  laboratory  in  the  preparation  of  hydro- 
sulphuric  acid  : 

FeS  +  2  HC1  =  FeCl2  +  H2S. 

Iron  disulphide  (pyrites]  (FeS2).  This  substance  bears 
the  same  relation  to  ferrous  sulphide  that  hydrogen  dioxide 
does  to  water.  It  occurs  abundantly  in  nature  in  the  form 
of  brass-yellow  cubical  crystals  and  in  compact  masses. 
Sometimes  the  name  "  fool's  gold  "  is  applied  to  it  from  its 
superficial  resemblance  to  the  precious  metal.  It  is  used  in 
very  large  quantities  as  a  source  of  sulphur  dioxide  in  the 
manufacture  of  sulphuric  acid,  since  it  burns  readily  in  the 
air,  forming  ferric  oxide  and  sulphur  dioxide : 

2  FeS2  +  1 1  O  =  Fe2O3  +  4  SO2. 

Ferrous  carbonate  (FeCO3).  This  compound  occurs  in 
nature  as  siderite,  and  is  a  valuable  ore.  It  will  dissolve 
to  some  extent  in  water  containing  carbon  dioxide,  just  as 
will  calcium  carbonate,  and  waters  containing  it  are  called 
chalybeate  waters.  These  chalybeate  waters  are  supposed 
to  possess  certain  medicinal  virtues  and  form  an  important 
class  of  mineral  waters. 

Ferric  salts.  Ferric  salts  are  usually  obtained  by  treating 
an  acidified  solution  of  a  ferrous  salt  with  an  oxidizing 
agent : 

2  FeCl2  +  2  HC1  +  O  =  2  FeCt3  +  H2O, 
2  FeSO4  +  H2SO4  +  O  =  Fe2(SO4)3  +  H2O. 

They  are  usually  yellow  or  violet  in  color,  are  quite  soluble, 
and  as  a  rule  do  not  crystallize  well.  Heated  with  water  in 
the  absence  of  free  acid,  they  hydrolyze  even  more  readily 
than  the  salts  of  aluminium.  The  most  familiar  ferric  salts 
are  the  chloride  and  the  sulphate. 


352     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Ferric  chloride  (FeCl3).  This  salt  can  be  obtained  most 
conveniently  by  dissolving  iron  in  hydrochloric  acid  and 
then  passing  chlorine  into  the  solution  : 

Fe  +  2  HC1  =  FeCl2  +  2  H, 
FeCl2  +  Cl  =  FeCl3. 

When  the  pure  salt  is  heated  with  water  it  is  partly  hydro- 
lyzed  :         p^  +  ^  H^Q  _^  Fe(OH)3  +  3  HC1. 


This  is  a  reversible  reaction,  however,  and  hydrolysis  can 
therefore  be  prevented  by  first  adding  a  considerable 
amount  of  the  soluble  product  of  the  reaction,  namely, 
hydrochloric  acid. 

Ferric  sulphate  (Fe2(SO4)3).  This  compound  can  be 
made  by  treating  an  acid  solution  of  green  vitriol  with 
an  oxidizing  agent.  It  is  difficult  to  crystallize  and  hard 
to  obtain  in  pure  condition.  When  an  alkali  sulphate  in 
proper  quantity  is  added  to  ferric  sulphate  in  solution 
an  iron  alum  is  formed,  and  is  easily  obtained  in  large 
crystals.  The  best  known  iron  alums  have  the  formulas 
KFe(SO4)2  •  12  H2O  and  NH4Fe(SO4)2  •  12  H2O.  They  are 
commonly  used  when  a  pure  ferric  salt  is  required. 

Ferric  hydroxide  (Fe(OH)8).  When  solutions  of  ferric 
salts  are  treated  with  ammonium  hydroxide,  ferric  hydroxide 
is  formed  as  a  rusty-red  precipitate,  insoluble  in  water. 

Iron  cyanides.  A  large  number  of  complex  cyanides  con- 
taining iron  are  known,  the  most  important  being  potassium 
ferrocyanide,  or  yellow  prussiate  of  potash  (K4FeC6N6), 
and  potassium  ferricyanide,  or  red  prussiate  of  potash 
(K3FeC6N6).  These  compounds  are  the  potassium  salts 
of  the  complex  acids  of  the  formulas  H4FeC6N6  and 
HoFeCJSL. 


THE  IRON  FAMILY  353 

Oxidation  of  ferrous  salts.  It  has  just  been  seen  that 
when  a  ferrous  salt  is  treated  with  an  oxidizing  agent  in 
the  presence  of  a  free  acid  a  ferric  salt  is  formed  : 

2  FeSO4  +  H2SO4  +  O  =  Fe2(SO4)3  +  H2O. 

In  this  reaction  oxygen  is  used  up,  and  the  valence  of  the 
iron  is  changed  from  2  to  3.  The  same  equation  may  be 
written 

2  Fe++,  2  SQ4-  +  2  H+,  SO4-  +  O  =  2  Fe+++,  3  SO4-+  H2O. 

Hydrogen  ions  have  been  oxidized  to  water,  while  the 
charge  of  each  iron  ion  has  been  increased  from  2  to  3. 

In  a  similar  way  the  conversion  of  ferrous  chloride  into 
ferric  chloride  may  be  written 

Fe++,  2  Cl-  +  Cl  =  Fe+++,  +  3  Cr. 

Here  again  the  valence  of  the  iron  and  the  charge  on  the 
iron  ion  has  been  increased  from  2  to  3,  though  no  oxygen 
has  entered  into  the  reaction.  As  a  rule,  however,  changes 
of  this  kind  are  brought  about  by  the  use  of  an  oxidizing 
agent,  and  are  called  oxidations. 

The  term  "  oxidation  "  is  applied  to  all  reactions  in  which 
the  valence  of  the  metal  of  a  compound  is  increased,  or, 
in  other  words,  to  all  reactions  in  which  the  charge  of  a 
cation  is  increased. 

Reduction  of  ferric  salts.  The  changes  which  take  place 
when  a  ferric  salt  is  converted  into  a  ferrous  salt  are  the 
reverse  of  the  ones  just  described.  This  is  seen  in  the 
equation  p^  +  R  =  p^  +  HQ 

In  this  reaction  the  valence  of  the  iron  has  been  changed 
from  3  to  2.  The  same  equation  may  be  written 

Fe+++,  3  Cl-  +  H  =  Fe++,  +  H+  +  3  Cl' 


354     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

It  will  be  seen  that  the  charge  of  the  iron  ions  has  been 
diminished  from  3  to  2.  Since  these  changes  are  the  re- 
verse of  the  oxidation  changes  just  considered,  they  are 
called  reduction  reactions.  The  term  "  reduction  "  is  applied 
to  all  processes  in  which  the  valence  of  the  metal  of  a  com- 
pound is  diminished,  or,  in  other  words,  to  all  processes  in 
which  the  charge  on  the  cations  is  diminished. 


NICKEL  AND  COBALT 

These  elements  occur  sparingly  in  nature,  usually  com- 
bined with  arsenic  or  with  arsenic  and  sulphur.  Both  ele- 
ments have  been  found  in  the  free  state  in  meteorites.  Like 
iron  they  form  two  series  of  compounds,  but  the  salts  cor- 
responding to  the  ferrous  salts  are  the  most  common, 
the  ones  corresponding  to  the  ferric  salts  being  difficult 
to  obtain.  Thus  we  have  the  chlorides  NiCl2  •  6  H2O 
and  CoCl2  •  6  H2O  ;  the  sulphates  NiSO4  •  7  H2O  and 
CoSO4-;H2O;  the  nitrates  Ni(NO3)2  •  6  H2O  and 
Co(NO3)2-6H2O. 

Nickel  is  largely  used  as  an  alloy  with  other  metals. 
Alloyed  with  copper  it  forms  coin  metal  from  which  five- 
cent  pieces  are  made,  with  copper  and  zinc  it  forms  Ger- 
man silver,  and  when  added  to  steel  in  small  quantities 
nickel  steel  is  formed  which  is  much  superior  to  common 
steel  for  certain  purposes.  When  deposited  by  electrolysis 
upon  the  surface  of  other  metals  such  as  iron,  it  forms  a 
covering  which  will  take  a  high  polish  and  protects  the 
metal  from  rust,  nickel  not  being  acted  upon  by  moist  air. 
Salts  of  nickel  are  usually  green. 

Compounds  of  cobalt  fused  with  glass  give  it  an  intensely 
blue  color.  In  powdered  form  such  glass  is  sometimes  used 


THE  IRON  FAMILY 


355 


as  a  pigment  called  smalt.  Cobalt  salts,  which  contain 
water  of  crystallization,  are  usually  cherry  red  in  color; 
when  dehydrated  they  become  blue. 


EXERCISES 

1.  In  the  manufacture  of  cast  iron,  why  is  the  air  heated  before 
being  forced  into  the  furnace? 

2.  Write  the  equations  showing  how  each  of  the  following  com- 
pounds of  iron  could  be  obtained  from  the  metal   itself  :   ferrous 
chloride,  ferrous  hydroxide,  ferrous  sulphate,  ferrous  sulphide,  fer- 
rous carbonate,  ferric  chloride,  ferric  sulphate,  ferric  hydroxide. 

3.  Account  for  the  fact  that  a  solution  of  sodium  carbonate,  when 
added  to  a  solution  of  a  ferric  salt,  precipitates  an  hydroxide  and  not 
a  carbonate. 

4.  Calculate  the  percentage  of  iron  in  each  of  the  common  iron  ores. 

5.  One  ton  of  steel  prepared  by  the  Bessemer  process  is  found 
by  analysis  to  contain  0.2%  carbon.    What  is  the  minimum  weight  of 
carbon  which  must  be  added  in  order  that  the  steel  may  be  made  to 
take  a  temper  ? 


CHAPTER   XXVIII 
COPPER,  MERCURY,  AND  SILVER 


SYMBOL 

ATOMIC 
WEIGHT 

DENSITY 

MELTING 
POINT 

FORMULAS  OF  OXIDES 

"  ous  " 

"ic  " 

Copper     .     . 

Cu 

63.6 

8.89 

1084° 

Cu2O 

CuO 

Mercury   .     . 

Hg 

200.00 

I3-596 

-39-5° 

Hg20 

HgO 

Silver  .     .     . 

Ag 

107.93 

I0.5 

960° 

Ag20 

AgO 

The  family.  By  referring  to  the  periodic  arrangement 
of  the  elements  (page  168),  it  will  be  seen  that  mercury  is 
not  included  in  the  same  family  with  copper  and  silver. 
Since  the  metallurgy  of  the  three  elements  is  so  similar, 
however,  and  since  they  resemble  each  other  so  closely  in 
chemical  properties,  it  is  convenient  to  class  them  together 
for  study. 

1.  Occurrence.    The  three  elements  occur  in  nature  to 
some  extent  in  the  free  state,  but  are  usually  found  as 
sulphides.    Their  ores  are  easy  to  reduce. 

2.  Properties.    They   are   heavy   metals   of    high   luster 
and  are  especially  good  conductors  of  heat  and  electricity. 
They  are  not  very  active  chemically.    Neither  hydrochloric 
nor  dilute  sulphuric  acid  has  any  appreciable  action  upon 
them.    Concentrated  sulphuric  acid  attacks  all  three,  form- 
ing metallic  sulphates  and  evolving  sulphur  dioxide,  while 
nitric  acid,  both  dilute  and  concentrated,  converts  them 
into  nitrates  with  the  evolution  of  oxides  of  nitrogen. 

356 


COPPER,  MERCURY,  AND  SILVER  357 

3.  Two  series  of  salts.  Copper  and  mercury  form  oxides 
of  the  types  M2O  and  MO,  as  well  as  two  series  of  salts. 
In  one  series  the  metals  are  univalent  and  the  salts  have 
formulas  like  those  of  the  sodium  salts.  They  are  called 
cuprous  and  mercurous  salts.  In  the  other  series  the  metals 
are  divalent  and  resemble  magnesium  salts  in  formulas. 
These  are  called  cupric  and  mercuric  salts.  Silver  forms 
only  one  series  of  salts,  being  always  a  univalent  metal. 

COPPER 

Occurrence.  The  element  copper  has  been  used  for 
various  purposes  since  the  earliest  days  of  history.  It  is 
often  found  in  the  metallic  state  in  nature,  large  masses  of 
it  occurring  pure  in  the  Lake  Superior  region  and  in  other 
places  to  a  smaller  extent.  The  most  valuable  ores  are  the 
following : 

Cuprite Cu2O. 

Chalcocite  . Cu2S. 

Chalcopyrite CuFeS2. 

Bornite Cu3FeS3. 

Malachite CuCO3  •  Cu(OH)2. 

Azurite 2  CuCO3  •  Cu(OH)2. 

Metallurgy  of  copper.  Ores  containing  little  or  no  sul- 
phur are  easy  to  reduce.  They  are  first  crushed  and  the 
earthy  impurities  washed  away.  The  concentrated  ore  is 
then  mixed  with  carbon  and  heated  in  a  furnace,  metallic 
copper  resulting  from  the  reduction  of  the  copper  oxide  by 
the  hot  carbon. 

Metallurgy  of  sulphide  ores.  Much  of  the  copper  of  commerce  is 
made  from  chalcopyrite  and  bornite,  and  these  ores  are  more  difficult 
to  work.  They  are  first  roasted  in  the  air,  by  which  treatment  much 
of  the  sulphur  is  burned  to  sulphur  dioxide.  The  roasted  ore  is  then 


358     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

melted  in  a  small  blast  furnace  or  in  an  open  one  like  a  puddling 
furnace.  In  melting,  part  of  the  iron  combines  with  silica  to  form 
a  slag  of  iron  silicate.  The  product,  called  crude  matte,  contains 
about  50%  copper  together  with  sulphur  and  iron.  Further  puri- 
fication is  commonly  carried  on  by  a  process  very  similar  to  the 
Bessemer  process  for  steel.  The  converter  is  lined  with  silica,  and 
a  charge  of  matte  from  the  melting  furnace,  together  with  sand,  is 
introduced,  and  air  is  blown  into  the  mass.  By  this  means  the  sul- 
phur is  practically  all  burned  out  by  the  air,  and  the  remaining  iron 
combines  with  silica  and  goes  off  as  slag.  The  copper  is  poured 
out  of  the  converter  and  molded  into  anode  plates  for  refining. 

Refining  of  copper.  Impure  copper  is  purified  by  elec- 
trolysis. A  large  plate  of  it,  serving  as  an  anode,  is  sus- 
pended in  a  tank  facing  a  thin  plate  of  pure  copper,  which 
is  the  cathode.  The  tank  is  filled  with  a  solution  of  copper 
sulphate  and  sulphuric  acid  to  serve  as  the  electrolyte.  A 
current  from  a  dynamo  passes  from  the  anode  to  the  cath- 
ode, and  the  copper,  dissolving  from  the  anode,  is  deposited 
upon  the  cathode  in  pure  form,  while  the  impurities  collect 
on  the  bottom  of  the  tank.  Electrolytic  copper  is  one  of 
the  purest  of  commercial  metals  and  is  very  nearly  pure 
copper. 

Recovery  of  gold  and  silver.  Gold  and  silver  are  often  present  in 
small  quantities  in  copper  ores,  and  in  electrolytic  refining  these 
metals  collect  in  the  muddy  deposit  on  the  bottom  of  the  tank.  The 
mud  is  carefully  worked  over  from  time  to  time  and  the  precious 
metals  extracted  from  it.  A  surprising  amount  of  gold  and  silver 
is  obtained  in  this  way. 

Properties  of  copper.  Copper  is  a  rather  heavy  metal  of 
density  8.9,  and  has  a  characteristic  reddish  color.  It  is 
rather  soft  and  is  very  malleable,  ductile,  and  flexible,  yet 
tough  and  strong;  it  melts  at  1084°.  As  a  conductor  of 
heat  and  electrical  energy  it  is  second  only  to  silver. 


COPPER,  MERCURY,  AND  SILVER  359 

Hydrochloric  acid,  dilute  sulphuric  acid,  and  fused  alkalis 
are  almost  without  action  upon  it ;  nitric  acid  and  hot,  con- 
centrated sulphuric  acid,  however,  readily  dissolve  it.  In 
moist  air  it  slowly  becomes  covered  with  a  thin  layer  of 
green  basic  carbonate  ;  heated  in  the  air  it  is  easily  oxidized 
to  black  copper  oxide  (CuO). 

Uses.  Copper  is  extensively  used  for  electrical  pur- 
poses, for  roofs  and  cornices,  for  sheathing  the  bottom  of 
ships,  and  for  making  alloys.  In  the  following  table  the 
composition  of  some  of  these  alloys  is  indicated  : 

COMPOSITION  OF  ALLOYS  OF  COPPER  IN  PERCENTAGES 

Aluminium  1 

±    .     copper  (90  to  97%),  aluminium  (3  to  10%). 

Brass.     .     .  .  copper  (63  to  73%),  zinc  (27  to  37%). 

Bronze    .     .  .  copper  (70  to  95%),  zinc  (i  to  25%),  tin  (i  to  18%). 

German  silver  .  copper  (56  to  60%),  zinc  (20%),  nickel  (20  to  25%). 

Gold  coin     .  .  copper  (10%),  gold  (90%). 

Gun  metal    .  .  copper  (90%),  tin  (10%). 

Nickel  coin.  .  copper  (75%),  nickel  (25%). 

Silver  coin  .  .  copper  (10%),  silver  (90%). 

Electrotyping.  Matter  is  often  printed  from  electrotype  plates 
which  are  prepared  as  follows.  The  matter  is  set  up  in  type  and 
wax  is  firmly  pressed  down  upon  the  face  of  it  until  a  clear  impres- 
sion is  obtained.  The  impressed  side  of  the  wax  is  coated  with 
graphite  and  the  impression  is  made  the  cathode  in  an  electrolytic 
cell  containing  a  copper  salt  in  solution.  When  connected  with  a 
current  the  copper  is  deposited  as  a  thin  sheet  upon  the  letters  in 
wax,  and  when  detached  is  a  perfect  copy  of  the  type,  the  under  part 
of  the  letters  being  hollow.  The  sheet  is  strengthened  by  pouring  on 
the  under  surface  a  suitable  amount  of  molten  metal  (commercial 
lead  is  used).  The  sheet  so  strengthened  is  then  used  in  printing. 

Two  series  of  copper  compounds.  Copper,  like  iron,  forms 
two  series  of  compounds  :  in  the  cuprous  compounds  it 
is  univalent ;  in  the  cupric  it  is  divalent.  The  cupric  salts 


360     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

are  much  the  more  common  of  the  two,  since  the  cuprous 
salts  pass  readily  into  cupric  by  oxidation. 

Cuprous  compounds.  The  most  important  cuprous  com- 
pound is  the  oxide  (Cu2O),  which  occurs  in  nature  as  ruby 
copper  or  cuprite.  It  is  a  bright  red  substance  and  can 
easily  be  prepared  by  heating  copper  to  a  high  tempera- 
ture in  a  limited  supply  of  air.  It  is  used  for  imparting 
a  ruby  color  to  glass. 

By  treating  cuprous  oxide  with  different  acids  a  number 
of  cuprous  salts  can  be  made.  Many  of  these  are  insoluble 
in  water,  the  chloride  (CuCl)  being  the  best  known.  When 
suspended  in  dilute  hydrochloric  acid  it  is  changed  into 
cupric  chloride,  the  oxygen  taking  part  in  the  reaction 
being  absorbed  from  the  air : 

2  CuCl  +  2  HC1  +  O=2  CuCl2  +  H2O. 

Cupric  compounds.  Cupric  salts  are  easily  made  by  dis- 
solving cupric  oxide  in  acids,  or,  when  insoluble,  by  pre- 
cipitation. Most  of  them  are  blue  or  green  in  color,  and 
the  soluble  ones  crystallize  well.  Since  they  are  so  much 
more  familiar  than  the  cuprous  salts,  they  are  frequently 
called  merely  copper  salts. 

Cupric  oxide  (CuO).  This  is  a  black  insoluble  substance 
obtained  by  heating  copper  in  excess  of  air,  or  by  igniting 
the  hydroxide  or  nitrate.  It  is  used  as  an  oxidizing  agent. 

Cupric  hydroxide  (Cu(OH)2).  The  hydroxide  prepared 
by  treating  a  solution  of  a  copper  salt  with  sodium  hydrox- 
ide is  a  light  blue  insoluble  substance  which  easily  loses 
water  and  changes  into  the  oxide.  Heat  applied  to  the 
liquid  containing  the  hydroxide  suspended  in  it  serves  to 
bring  about  the  reaction  represented  by  the  equation 

Cu(OH)2  =  CuO  +  H2O. 


COPPER,  MERCURY,  AND  SILVER  361 

Cupric  sulphate  (blue  vitriol]  (CuSO4- 5  H2O).  This 
substance,  called  blue  vitriol  or  bluestone,  is  obtained  as  a 
by-product  in  a  number  of  processes  and  is  produced  in 
very  large  quantities.  It  forms  large  blue  crystals,  which 
lose  water  when  heated  and  crumble  to  a  white  powder. 
The  salt  finds  many  uses,  especially  in  electro  typing  and 
in  making  electrical  batteries. 

Cupric  sulphide  (CuS).  The  insoluble  black  sulphide 
(CuS)  is  easily  prepared  by  the  action  of  hydrosulphuric 
acid  upon  a  solution  of  a  copper  salt : 

.    CuSO4  +  H2S  =  CuS  +  H2SO4. 
It  is  insoluble  in  water  and  dilute  acids. 


MERCURY 

Occurrence.  Mercury  occurs  in  nature  chiefly  as  the 
sulphide  (HgS)  called  cinnabar,  and  in  globules  of  metal 
inclosed  in  the  cinnabar.  The  mercury  mines  of  Spain 
have  long  been  famous,  California  being  the  next  largest 
producer. 

Metallurgy.  Mercury  is  a  volatile  metal  which  has  but 
little  affinity  for  oxygen.  Sulphur,  on  the  other  hand,  readily 
combines  with  oxygen.  These  facts  make  the  metallurgy 
of  mercury  very  simple.  The  crushed  ore,  mixed  with  a 
small  amount  of  carbon  to  reduce  any  oxide  or  sulphate 
that  might  be  formed,  is  roasted  in  a  current  of  air.  The 
sulphur  burns  to  sulphur  dioxide,  while  the  mercury  is  con- 
verted into  vapor  and  is  condensed  in  a  series  of  condens- 
ing vessels.  The  metal  is  purified  by  distillation. 

Properties.  Mercury  is  a  heavy  silvery  liquid  with  a 
density  of  13.596.  It  boils  at  357°  and  solidifies  at  — 39.5°. 


362     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Small  quantities  of  many  metals  dissolve  in  it,  forming 
liquid  alloys,  while  with  larger  quantities  it  forms  solid 
alloys.  The  alloys  of  mercury  are  called  amalgams. 

Toward  acids  mercury  conducts  itself  very  much  like 
copper;  it  is  easily  attacked  by  nitric  and  hot,  concentrated 
sulphuric  acids,  while  cold  sulphuric  and  hydrochloric  acids 
have  no  effect  on  it. 

Uses.  Mercury  is  extensively  used  in  the  construction 
of  scientific  instruments,  such  as  the  thermometer  and 
barometer,  and  as  a  liquid  over  which  to  collect  gases 
which  are  soluble  in  water.  The  readiness  with  which  it 
alloys  with  silver  and  gold  makes  it  very  useful  in  the 
extraction  of  these  elements. 

Compounds  of  mercury.  Like  copper,  mercury  forms  two 
series  of  compounds  :  the  mercurous,  of  which  mercurous 
chloride  (HgCl)  is  an  example ;  and  the  mercuric,  repre- 
sented by  mercuric  chloride  (HgCl2). 

Mercuric  oxide  (HgO).  Mercuric  oxide  can  be  obtained 
either  as  a  brick-red  or  as  a  yellow  substance.  When  mer- 
curic nitrate  is  heated  carefully  the  red  modification  is 
formed  in  accordance  with  the  equation 

Hg(N08)2  =  HgO  +  2N02  +  0. 

The  yellow  modification  is  prepared  by  adding  a  solution 
of  a  mercuric  salt  to  a  solution  of  sodium  or  potassium 
hydroxide  : 

Hg(NO3)2  +  2  NaOH  =  2  NaNO3  +  Hg(OH)2, 
Hg(OH)2  =  HgO  +  H20. 

When  heated  the  oxide  darkens  until  it  becomes  almost 
black ;  at  a  higher  temperature  it  decomposes  into  mer- 
cury and  oxygen.  It  was  by  this  reaction  that  oxygen  was 
discovered. 


COPPER,  MERCURY,  AND  SILVER  363 

Mercurous  chloride  (calomel]  (HgCl).  Being  insoluble, 
mercurous  chloride  is  precipitated  as  a  white  solid  when  a 
soluble  chloride  is  added  to  a  solution  of  mercurous  nitrate  : 

HgN03  +  NaCl  =  HgCl  +  NaNO3. 

Commercially  it  is  manufactured  by  heating  a  mixture  of 
mercuric  chloride  and  mercury.  When  exposed  to  the  light 
it  slowly  changes  into  mercuric  chloride  and  mercury : 

2  HgCl  =  HgCl2  +  Hg. 

It  is  therefore  protected  from  the  light  by  the  use  of 
colored  bottles.  It  is  used  in  medicine. 

Most  mercurous  salts  are  insoluble  in  water,  the  prin- 
cipal soluble  one  being  the  nitrate,  which  is  made  by  the 
action  of  cold,  dilute  nitric  acid  on  mercury. 

Mercuric  chloride  (corrosive  sublimate}  (HgCl2).  This 
substance  can  be  made  by  dissolving  mercuric  oxide  in 
hydrochloric  acid.  On  a  commercial  scale  it  is  made  by 
subliming  a  mixture  of  common  salt  and  mercuric  sulphate  : 

2  NaCl  +  HgSO4  =  HgCl2  +  Na2SO4. 

The  mercuric  chloride,  being  readily  volatile,  vaporizes 
and  is  condensed  again  in  cool  vessels.  Like  mercurous 
chloride  it  is  a  white  solid,  but  differs  from  it  in  that  it  is 
soluble  in  water.  It  is  extremely  poisonous  and  in  dilute 
solutions  is  used  as  an  antiseptic  in  dressing  wounds. 

Mercuric  sulphide  (HgS).  As  cinnabar  this  substance 
forms  the  chief  native  compound  of  mercury,  occurring  in 
red  crystalline  masses.  By  passing  hydrosulphuric  acid 
into  a  solution  of  a  mercuric  salt  it  is  precipitated  as  a 
black  powder,  insoluble  in  water  and  acids.  By  other 
means  it  can  be  prepared  as  a  brilliant  red  powder  known 
as  vermilion,  which  is  used  as  a  pigment  in  fine  paints. 


364     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

The  iodides  of  mercury.  If  a  solution  of  potassium  iodide  is  added 
to  solutions  of  a  mercurous  and  a  mercuric  salt  respectively,  the 
corresponding  iodides  are  precipitated.  Mercuric  iodide  is  the  more 
important  of  the  two,  and  as  prepared  above  is  a  red  powder  which 
changes  to  yellow  on  heating  to  1 50°.  The  yellow  form  on  cooling 
changes  back  again  to  the  red  form,  or  may  be  made  to  do  so  by 
rubbing  it  with  a  knife  blade  or  some  other  hard  object. 

SILVER 

Occurrence.  Silver  is  found  in  small  quantities  in  the 
uncombined  state ;  usually,  however,  it  occurs  in  com- 
bination with  sulphur,  either  as  the  sulphide  (Ag2S)  or  as  a 
small  constituent  of  other  sulphides,  especially  those  of  lead 
and  copper.  It  is  also  found  alloyed  with  gold. 

Metallurgy.  Parkess  process.  Silver  is  usually  smelted 
in  connection  with  lead.  The  ores  are  worked  over  together, 
as  described  under  lead,  and  the  lead  and  silver  obtained 
as  an  alloy,  the  silver  being  present  in  small  quantity.  The 
alloy  is  melted  and  metallic  zinc  is  stirred  in.  Zinc  will 
alloy  with  silver  but  not  with  lead,  and  it  is  found  that 
the  silver  leaves  the  lead  and,  in  the  form  of  an  alloy  with 
zinc,  forms  as  a  crust  upon  the  lead  and  is  skimmed  off. 
This  crust,  which,  of  course,  contains  lead  adhering  to  it,  is 
partially  melted  and  the  most  of  the  lead  drained  off.  The 
zinc  is  removed  by  distillation,  and  the  residue  is  melted 
on  an  open  hearth  in  a  current  of  air ;  by  this  means  the 
zinc  and  lead  remaining  with  the  silver  are  changed  into 
oxides  and  the  silver  remains  behind  unaltered. 

Amalgamation  process.  In  some  localities  the  old  amalgamation 
process  is  used.  The  silver  ore  is  treated  with  common  salt  and  fer- 
rous compounds,  which  process  converts  the  silver  first  into  chloride 
and  then  into  metallic  silver.  Mercury  is  then  added  and  thoroughly 
mixed  with  the  mass,  forming  an  amalgam  with  the  silver.  After 


COPPER,  MERCURY,  AND  SILVER  365 

some  days  the  earthy  materials  are  washed  away  and  the  heavier 
amalgam  is  recovered.  The  mercury  is  distilled  off  and  the  silver 
left  in  impure  form. 

Refining  silver.  The  silver  obtained  by  either  of  the 
above  processes  may  still  contain  copper,  gold,  and  iron, 
and  is  refined  by  "parting"  with  sulphuric  acid.  The 
metal  is  heated  with  strong  sulphuric  acid  which  dissolves 
the  silver,  copper,  and  iron  present,  but  not  the  gold.  In 
the  solution  of  silver  sulphate  so  obtained  copper  plates 
are  suspended,  upon  which  the  pure  silver  precipitates,  the 
copper  going  into  solution  as  sulphate,  as  shown  in  the 
equation 

Ag2SO4  +  Cu  =  2  Ag  +  CuSO4. 

The  solution  obtained  as  a  by-product  in  this  process  fur- 
nishes most  of  the  blue  vitriol  of  commerce.  Silver  is  also 
refined  by  electrolytic  methods  similar  to  those  used  in 
refining  copper. 

Properties  of  silver.  Silver  is  a  heavy,  rather  soft,  white 
metal,  very  ductile  and  malleable  and  capable  of  taking  a 
high  polish.  It  surpasses  all  other  metals  as  a  conductor 
of  heat  and  electricity,  but  is  too  costly  to'  find  extensive 
use  for  such  purposes.  It  melts  at  a  little  lower  tempera- 
ture than  copper  (961°).  It  alloys  readily  with  other  heavy 
metals,  and  when  it  is  to  be  used  for  coinage  a  small 
amount  of  copper  —  from  8  to  10%  —  is  nearly  always 
melted  with  it  to  give  it  hardness. 

It  is  not  acted  upon  by  water  or  air,  but  is  quickly  tar- 
nished when  in  contact  with  sulphur  compounds,  turning 
quite  black  in  time.  Hydrochloric  acid  and  fused  alkalis 
do  not  act  upon  it,  but  nitric  acid  and  hot,  concentrated 
sulphuric  acid  dissolve  it  with  ease. 


366     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


FIG. 


Electroplating.  Since  silver  is  not  acted  upon  by  water  or  air,  and 
has  a  pleasing  appearance,  it  is  used  to  coat  various  articles  made  of 
cheaper  metals.  Such  articles  are  said  to  be  silver  plated.  The  process 
by  which  this  is  done  is  called  electroplating.  It  is  carried  on  as  fol- 
lows :  The  object  to  be  plated  (such  as  a  spoon)  is  attached  to  a  wire 
and  dipped  into  a  solution  of  a  silver  salt.  Electrical  connection  is 

made  in  such  a 
way  that  the  arti- 
cle to  be  plated 
serves  as  the  cath- 
ode,  while  the 
anode  is  made  up 
of  one  or  more 
plates  of  silver 
(Fig.  88,  A). 

When  a  current  is  passed  through  the  electrolyte  silver  dissolves 
from  the  anode  plate  and  deposits  on  the  cathode  in  the  form  of  a 
closely  adhering  layer.  By  making  the  proper  change  in  the  electrolyte 
and  anode  plate  objects  may  be  plated  with  gold  and  other  metals. 

Compounds  of  silver.  Silver  forms  two  oxides  but  only 
one  series  of  salts,  namely,  the  one  which  corresponds  to 
the  mercurous  and  cuprous  series. 

Silver  nitrate  (lunar  caustic}  (AgNO3).  This  salt  is 
easily  prepared  by  dissolving  silver  in  nitric  -acid  and 
evaporating  the  resulting  solution.  It  crystallizes  in  flat 
plates,  and  when  heated  carefully  can  be  melted  without 
decomposition.  When  cast  into  sticks  it  is  called  lunar 
caustic,  for  it  has  a  very  corrosive  action  on  flesh,  and  is 
sometimes  used  in  surgery  to  burn  away  abnormal  growths. 

The  alchemists  designated  the  metals  by  the  names  of  the  heav- 
enly bodies.  The  moon  (luna)  was  the  symbol  for  silver  ;  hence  the 
name  "  lunar  caustic." 

Silver  sulphide  (Ag2S).  This  occurs  in  nature  and  con- 
stitutes one  of  the  principal  ores  of  silver.  It  can  be 


COPPER,  MERCURY,  AND  SILVER  367 

obtained  in  the  form  of  a  black  solid  by  passing  hydro- 
sulphuric  acid  through  a  solution  of  silver  nitrate. 

Compounds  of  silver  with  the  halogens.  The  chloride, 
bromide,  and  iodide  of  silver  are  insoluble  in  water  and 
acids,  and  are  therefore  precipitated  by  bringing  together 
a  soluble  halogen  salt  with  silver  nitrate  : 

AgNO3  +  KC1  =  AgCl  +  KNO3. 

They  are  remarkable  for  the  fact  that  they  are  very  sensi- 
tive to  the  action  of  light,  undergoing  a  change  of  color 
and  chemical  composition  when  exposed  to  sunlight,  espe- 
cially if  in  contact  with  organic  matter  such  as  gelatin. 

Photography.  The  art  of  photography  is  based  on  the  fact  that 
the  halogen  compounds  of  silver  are  affected  by  the  light,  particularly 
in  the  presence  of  organic  matter.  From  a  chemical  standpoint  the 
processes  involved  may  be  described  under  two  heads  :  (i)  the  prepa- 
ration of  the  negative  ;  (2)  the  preparation  of  the  print.  ' 

i.  Preparation  of  the  negative.  The  plate  used  in  the  preparation 
of  the  negative  is  made  by  spreading  a  thin  layer  of  gelatin,  in  which 
silver  bromide  is  suspended  (silver  iodide  is  sometimes  added  also), 
over  a  glass  plate  or  celluloid  film  and  allowing  it  to  dry.  When  the 
plate  so  prepared  is  placed  in  a  camera  and  the  image  of  some  object 
is  focused  upon  it,  the  silver  salt  undergoes  a  change  which  is  propor- 
tional at  each  point  to  the  intensity  of  the  light  falling  upon  it.  In 
this  way  an  image  of  the  object  photographed  is  produced  upon  the 
plate,  which  is,  however,  invisible  and  is  therefore  called  "  latent." 
It  can  be  made  visible  by  the  process  of  developing. 

To  develop  the  image  the  exposed  plate  is  immersed  in  a  solution 
of  some  reducing  agent  called  the  developer.  The  developer  reduces 
that  portion  of  the  silver  salt  which  has  been  affected  by  the  light, 
depositing  it  in  the  form  of  black  metallic  silver  which  closely  adheres 
to  the  plate. 

The  unaffected  silver  salt,  upon  which  the  developer  has  no  action, 
must  now  be  removed  from  the  plate.  This  is  done  by  immersing 
the  plate  in  a  solution  of  sodium  thiosulphate  (hypo).  After  the  sil- 
ver salt  has  been  dissolved  off,  the  plate  is  washed  with  water  and 


368     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

dried.  The  plate  so  prepared  is  called  the  negative  because  it  is  a 
picture  of  the  object  photographed,  with  the  lights  exactly  reversed. 
This  is  called  fixing  the  negative. 

2.  Preparation  of  the  print.  The  print  is  made  from  paper 
which  is  prepared  in  the  same  way  as  the  negative  plate.  The  nega- 
tive is  placed  upon  this  paper  and  exposed  to  the  light  in  such  a  way 
that  the  light  must  pass  through  the  negative  before  striking  the 
paper.  If  the  paper  is  coated  with  silver  chloride,  a  visible  image  is 
produced,  in  which  case  a  developer  is  not  needed.  The  proofs  are 
made  in  this  way.  In  order  to  make  them  permanent  the  unchanged 
silver  chloride  must  be  dissolved  off  with  sodium  thiosulphate.  The 
print  is  then  toned  by  dipping  it  into  a  solution  of  gold  or  platinum 
salts.  The  silver  on  the  print  passes  into  solution,  while  the  gold  or 
platinum  takes  its  place.  These  metals  give  a  characteristic  color  or 
tone  to  the  print,  the  gold  making  it  reddish  brown,  while  the  platinum 
gives  it  a  steel-gray  tone.  If  a  silver  bromide  paper  is  used  in  making 
the  print,  a  latent  image  is  produced  which  must  be  developed  as  in 
the  case  of  the  negative  itself.  The  silver  bromide  is  much  more  sen- 
sitive than  the  chloride,  so  that  the  printing  can  be  done  in  artificial 
light.  Since- the  darkest  places  on  the  negative  cut  off  the  most  light, 
it  is  evident  that  the  lights  of  the  print  will  be  the  reverse  of  those  of 
the  negative,  and  will  therefore  correspond  to  those  of  the  object  pho- 
tographed. The  print  is  therefore  called  the  positive. 

EXERCISES 

1.  Account  for  the  fact  that  copper  has  been  used  for  so  long  a 
time. 

2.  Write  equations  for  the  action  of  concentrated  sulphuric  and 
nitric  acids  upon  the  metals  of  this  family. 

3.  How  would  you  account  for  the  fact  that  normal  copper  sul- 
phate is  slightly  acid  to  litmus  ? 

4.  Contrast  the  action  of  heat  on  cupric  nitrate   and   mercuric 
nitrate. 

5.  State  reasons  why  mercury  is  adapted  for  use  in  thermometers 
and  barometers. 

6.  How  could  you  distinguish  between  mercurous  chloride  and 
mercuric  chloride? 

7.  Write  equations  for  the  preparation  of  mercuric  and  mercurous 
iodides. 


COPPER,  MERCURY,  AND  SILVER  369 

8.  How  would  you  account  for  the  fact  that  solutions  of  the  differ- 
ent salts  of  a  metal  usually  have  the  same  color  ? 

9.  Crude   silver   usually  contains   iron   and  lead.    What  would 
become  of  these  metals  in  refining  by  parting  with  sulphuric  acid  ? 

10.  In  the  amalgamation  process  for  extracting  silver,  how  does 
ferrous  chloride  convert  silver  chloride  into  silver  ?    Write  equation. 
Why  is  the  silver  sulphide  first  changed  into  silver  chloride? 

11.  What  impurities  would  you  expect  to  find  in  the  copper  sul- 
phate prepared  from  the  refining  of  silver? 

12.  How  could  you  prepare  pure  silver  chloride  from  a  silver  coin  ? 

13.  Mercuric  nitrate  and  silver  nitrate  are  both  white  solids  solu- 
ble in  water.    How  could  you  distinguish  between  them? 

14.  Account  for  the  fact  that  sulphur  waters  turn  a  silver  coin 
black  ;  also  for  the  fact  that  a  silver  spoon  is  blackened  by  foods 
(eggs,  for  example)  containing  sulphur. 

15.  When  a  solution  of  silver  nitrate  is  added  to  a  solution  of 
potassium  chlorate  no  precipitate  forms.    How  do  you  account  for 
the  fact  that  a  precipitate  of  silver  chloride  is  not  formed? 


CHAPTER    XXIX 
TIN  AND  LEAD 


SYMBOL 

ATOMIC 
WEIGHT 

DENSITY 

MELTING 
POINT 

COMMON  OXIDES 

Tin.     ... 

Sn 

IIQ.O 

7-35 

235° 

SnO                   SnO2 

Lead    .     .     . 

Pb 

206.9 

11.38 

327° 

PbO    Pb3O4    PbO2 

The  family.  Tin  and  lead,  together  with  silicon  and 
germanium,  form  a  family  in  Group  IV  of  the  periodic 
table.  Silicon  has  been  discussed  along  with  the  non- 
metals,  while  germanium,  on  account  of  its  rarity,  needs 
only  to  be  mentioned. 

The  other  family  of  Group  IV  includes  carbon,  already 
described,  and  a  number  of  rare  elements. 


TIN 

Occurrence.  Tin  is  found  in  nature  chiefly  as  the  oxide 
(SnO2),  called  cassiterite  or  tinstone.  The  most  famous 
mines  are  those  of  Cornwall  in  England,  and  of  the  Malay 
Peninsula  and  East  India  Islands ;  in  small  amounts  tin- 
stone is  found  in  many  other  localities. 

Metallurgy.  The  metallurgy  of  tin  is  very  simple.  The 
ore,  separated  as  far  as  possible  from  earthy  materials, 
is  mixed  with  carbon  and  heated  in  a  furnace,  the  reduc- 
tion taking  place  readily.  The  equation  is 

SnO   +  C  =  Sn  +  CO. 


TIN  AND  LEAD 


371 


The  metal  is  often  purified  by  carefully  heating  it  until  it 
is  partly  melted  ;  the  pure  tin  melts  first  and  can  be 
drained  away  from  the  impurities. 

Properties.  Pure  tin,  called  block  tin,  is  a  soft  white 
metal  with  a  silver-like  appearance  and  luster;  it  melts 
readily  (235°)  and  is  somewhat  lighter  than  copper,  having 
a  density  of  7.3.  It  is  quite  malleable  and  can  be  rolled 
out  into  very  thin  sheets,  forming  tin  foil  ;  most  tin  foil, 
however,  contains  a  good  deal  of  lead. 

Under  ordinary  conditions  it  is  quite  unchanged  by  air 
or  moisture,  but  at  a  high  temperature  it  burns  in  air, 
forming  the  oxide  SnO2.  Dilute  acids  have  no  effect  upon 
it,  but  concentrated  acids  attack  it  readily.  Concentrated 
hydrochloric  acid  changes  it  into  the  chloride 


With  sulphuric  acid  tin  sulphate  and  sulphur  dioxide  are 
formed  :  Sn  +  2  H2SO4  =  SnSO4  +  SO2  .+  2  H2O. 

Concentrated  nitric  acid  oxidizes  it,  forming  a  white  in- 
soluble compound  of  the  formula  H2SnO3,  called  meta- 
stannic  acid  : 

3  Sn  +  4  HNO3  +  H2O  =  3  H2SnO3  +  4  NO. 

Uses  of  tin.  A  great  deal  of  tin  is  made  into  tin  plate 
by  dipping  thin  steel  sheets  into  the  melted  metal.  Owing 
to  the  way  in  which  tin  resists  the  action  of  air  and  dilute 
acids,  tin  plate  is  used  in  many  ways,  such  as  in  roofing, 
and  in  the  manufacture  of  tin  cans,  cooking  vessels,  and 
similar  articles. 

Many  useful  alloys  contain  tin,  some  of  which  have 
been  mentioned  in  connection  with  copper.  When  tin  is 
alloyed  with  other  metals  of  low  melting  point,  soft,  easily 


372     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

melted  alloys  are  formed  which  are  used  for  friction  bear- 
ings in  machinery  ;  tin,  antimony,  lead,  and  bismuth  are 
the  chief  constituents  of  these  alloys.  Pewter  and  soft 
solder  are  alloys  of  tin  and  lead. 

Compounds  of  tin.  Tin  forms  two  series  of  compounds: 
the  stannous,  in  which  the  tin  is  divalent,  illustrated  in  the 
compounds  SnO,  SnS,  SnCl2 ;  the  stannic,  in  which  it 
is  tetravelent  as  shown  in  the  compounds  SnO2,  SnS2. 
There  is  also  an  acid,  H2SnO3,  called  stannic  acid,  which 
forms  a  series  of  salts  called  stannates.  While  this  acid 
has  the  same  composition  as  metastannic  acid,  the  two  are 
quite  different  in  their  chemical  properties.  This  differ- 
ence is  probably  due  to  the  different  arrangement  of  the 
atoms  in  the  molecules  of  the  two  substances.  Only  a  few 
compounds  of  tin  need  be  mentioned. 

Stannic  oxide  (SnO2).  Stannic  oxide  is  of  interest,  since 
it  is  the  chief  compound  of  tin  found  in  nature.  It  is 
sometimes  found  in  good-sized  crystals,  but  as  prepared  in 
the  laboratory  is  a  white  powder.  When  fused  with  potas- 
sium hydroxide  it  forms  potassium  stannate,  acting  very 
much  like  silicon  dioxide  : 

SnO2  +  2  KOH  =  K2SnO3  +  H2O. 

Chlorides  of  tin.  Stannous  chloride  is  prepared  by  dis- 
solving tin  in  concentrated  hydrochloric  acid  and  evap- 
orating the  solution  to  crystallization.  The  crystals  which 
are  obtained  have  the  composition  SnCl2'2  H2O,  and  are 
known  as  tin  crystals.  By  treating  a  solution  of  stannous 
chloride  with  aqua  regia,  stannic  chloride  is  formed : 

SnCl2  +  2  Cl  =  SnCl4. 

The  salt  which  crystallizes  from  such  a  solution  has  the 
composition  SnCl4  •  5  H2O,  and  is  known  commercially  as 


TIN  AND  LEAD 


373 


oxymuriate  of  tin.  If  metallic  tin  is  heated  in  a  current  of 
dry  chlorine,  the  anhydrous  chloride  (SnCl4)  is  obtained  as 
a  heavy  colorless  liquid  which  fumes  strongly  on  exposure 
to  air. 

The  ease  with  which  stannous  chloride  takes  up  chlorine 
to  form  stannic  chloride  makes  it  a  good  reducing  agent 
in  many  reactions,  changing  the  higher  chlorides  of  metals 
to  lower  ones.  Thus  mercuric  chloride  is  changed  into 
mercurous  chloride  : 

SnCl2  +  2  HgCl2  =  SnCl4  +  2  HgCL 

If  the  stannous  chloride  is  in  excess,  the  reaction  may  go 
further,  producing  metallic  mercury  : 

SnCl2  +  2  HgCl  =  SnCl4  +  2  Hg. 

Ferric  chloride  is  in  like  manner  reduced  to  ferrous  chlo- 
nde  :  SnCl2  +  2  FeCl3  =  SnCl4  +  2  FeCl2. 

The  chlorides  of  tin,  as  well  as  the  alkali  stannates,  are 
much  used  as  mordants  in  dyeing  processes.  The  hydrox- 
ides of  tin  and  free  stannic  acid,  which  are  easily  liberated 
from  these  compounds,  possess  in  very  marked  degree  the 
power  of  fixing  dyes  upon  fibers,  as  explained  under  alu- 
minium. 

LEAD 

Occurrence.  Lead  is  found  in  nature  chiefly  as  the  sul- 
phide (PbS),  called  galena;  to  a  much  smaller  extent  it 
occurs  as  carbonate,  sulphate,  c'hromate,  and  in  a  few 
other  forms.  Practically  all  the  lead  of  commerce  is  made 
from  galena,  two  general  methods  of  metallurgy  being  in  use. 

Metallurgy,  i .  The  sulphide  is  melted  with  scrap  iron, 
when  iron  sulphide  and  metallic  lead  are  formed ;  the 


374     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

liquid  lead,  being  the  heavier,  sinks  to  the  bottom  of  the 
vessel  and  can  be  drawn  off : 

PbS  +  Fe  =  Pb  +  FeS. 

2.  The  sulphide  is  roasted  in  the  air  until  a  part  of  it 
has  been  changed  into  oxide  and  sulphate.  The  air  is  then 
shut  off  and  the  heating  continued,  the  reactions  indicated 
in  the  following  equations  taking  place  : 

2  PbO  +  PbS  =  3  Pb  +  SO2, 
PbSO4  +  PbS  =  2  Pb  +  2  SO2. 

The  lead  so  prepared  usually  contains  small  amounts 
of  silver,  arsenic,  antimony,  copper,  and  other  metals. 
The  silver  is  removed  by  Parkes's  method,  as  described 
under  silver,  and  the  other  metals  in  various  ways.  The 
lead  of  commerce  is  one  of  the  purest  commercial  metals, 
containing  as  a  rule  only  a  few  tenths  per  cent  of  im- 
purities. 

Properties.  Lead  is  a  heavy  metal  (den.  =  11.33)  which 
has  a  brilliant  silvery  luster  on  a  freshly  cut  surface,  but 
which  soon  tarnishes  to  a  dull  blue-gray  color.  It  is  soft, 
easily  fused  (melting  at  327°),  and  quite  malleable,  but  has 
little  toughness  or  strength. 

It  is  not  acted  upon  to  any  great  extent  by  the  oxygen 
of  the  air  under  ordinary  conditions,  but  is  changed  into 
oxide  at  a  high  temperature.  With  the  exception  of  hydro- 
chloric and  sulphuric  acids,  most  acids,  even  very  weak 
ones,  act  upon  it,  forming  soluble  lead  salts.  Hot,  concen- 
trated hydrochloric  and  sulphuric  acids  also  attack  it  to 
a  slight  extent. 

Uses.  Lead  is  employed  in  the  manufacture  of  lead 
pipes  and  in  large  storage  batteries.  In  the  form  of  sheet 
lead  it  is  used  in  lining  the  chambers  of  sulphuric  acid 


TIN  AND  LEAD  375 

works  and  in  the  preparation  of  paint  pigments.  Some 
alloys  of  lead,  such  as  solder  and  pewter  (lead  and  tin), 
shot  (lead  and  arsenic),  and  soft  bearing  metals,  are 
widely  used.  Type  metal  consists  of  lead,  antimony,  and 
sometimes  tin.  Compounds  of  lead  form  several  impor- 
tant pigments. 

Compounds  of  lead.  In  nearly  all  its  compounds  lead 
has  a  valence  of  2,  but  a  few  corresponding  to  stannic 
compounds  have  a  valence  of  4. 

Lead  oxides.  Lead  forms  a  number  of  oxides,  the  most 
important  of  which  are  litharge,  red  lead  or  minium,  and 
lead  peroxide. 

1.  Litharge  (PbO).    This  oxide  forms  when  lead  is  oxi- 
dized at  a  rather  low  temperature,  and   is   obtained   as  a 
by-product  in  silver  refining.    It  is  a  pale  yellow  powder, 
and  has  a  number  of  commercial  uses.    It  is  easily  soluble 
in  nitric  acid  : 

PbO  +  2  HNO3  -  Pb(NO3)2  +  H2O. 

2.  Red  lead,  or  minium  (Pb3O4).    Minium  is  prepared  by 
heating  lead  (or  litharge)  to  a  high  temperature  in  the  air. 
It  is  a  heavy  powder  of  a  beautiful  red  color,  and  is  much 
used  as  a  pigment. 

3.  Lead  peroxide  (PbO2).    This  is  left  as  a  residue  when 
minium  is  heated  with  nitric  acid  : 

Pb3O4  4-  4  HNO3  =  2  Pb(NO3)2  +  PbO2  -f  2  H2O. 

It  is  a  brown  powder  which  easily  gives  up  a  part  of  its 
oxygen  and,  like  manganese  dioxide  and  barium  dioxide,  is 
a  good  oxidizing  agent. 

Soluble  salts  of  lead.  The  soluble  salts  of  lead  can  be 
made  by  dissolving  litharge  in  acids.  Lead  acetate 
(Pb(C2H3O2)2  •  3  H2O),  called  sugar  of  lead,  and  lead  nitrate 


376     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


(Pb(NO3)2)  are  the  most  familiar  examples.    They  are  white 
crystalline  solids  and  are  poisonous  in  character. 

Insoluble  salts  of  lead;  lead  carbonate.  While  the  nor- 
mal carbonate  of  lead  (PbCO3)  is  found  to  some  extent 
in  nature  and  can  be  prepared  in  the  laboratory,  basic 
carbonates  of  varying  pomposition  are  much  more  easy  to 
obtain.  One  of  the  simplest  of  these  has  the  composition 
2.  PbCO3  •  Pb(OH)2.  A  mixture  of  such  carbonates  is  called 
white  lead.  This  is  prepared  on  a  large  scale  as  a  paint 
pigment  and  as  a  body  for  paints  which  are  to  be  colored ' 
with  other  substances. 

White  lead.  White  lead  is  an  amorphous  white  substance  which, 
when  mixed  with  oil,  has  great  covering  power,  that  is,  it  spreads  out 
in  an  even  waxy  film,  free  from,  streaks  and  lumps,  and  covers  the 
entire  surface  upon  which  it  is  spread.  Its  disadvantage  as  a  pigment 
lies  in  the  fact  that  it  gradually  blackens  when  exposed  to  sulphur 
compounds,  which  are  often  present  in  the 
air,  forming  black  lead  sulphide  (PbS). 

Technical  preparation  of  white  lead.  Dif- 
ferent methods  are  used  in  the  preparation 
of  white  lead,  but  the  old  one  known  as 
the  Dutch  process  is  still  the  principal  one 
employed.  In  this  process,  earthenware 
pots  about  ten  inches  high  and  of  the  shape 
shown  in  Fig.  89  are  used.  In  the  bottom 
A  is  placed  a  3%  solution  of  acetic  acid 
(vinegar  answers  the  purpose  very  well). 
The  space  above  this  is  filled  with  thin, 
perforated,  circular  pieces  of  lead,  sup- 
ported by  the  flange  B  of  the  pot.  These 
pots  are  placed  close  together  on  a  bed  of  tan  bark  on  the  floor  of  a 
room  known  as  the  corroding  room.  They  are  covered  over  with 
boards,  upon  which  tan  bark  is  placed,  and  another  row  of  pots  is 
placed  on  this.  In  this  way  the  room  is  filled.  The  white  lead  is 
formed  by  the  fumes  of  the  acetic  acid,  together  with  the  carbon 
dioxide  set  free  in  the  fermentation  of  the  tan  bark  acting  on  the 
lead.  About  three  months  are  required  to  complete  the  process. 


FIG. 89 


TIN  AND  LEAD  377 

Lead  sulphide  (PbS).  In  nature  this  compound  occurs 
in  highly  crystalline  condition,  the  crystals  having  much 
the  same  luster  as  pure  lead.  It  is  readily  prepared  in  the 
laboratory  as  a  black  precipitate,  by  the  action  of  hydro- 
sulphuric  acid  upon  soluble  lead  salts  : 

Pb(NO3)2  +  H2S  =  PbS  +  2  HNO3. 

It  is  insoluble  both  in  water  and  in  dilute  acids. 

Other  insoluble  salts.  Lead  chromate  (PbCrO4)  is  a 
yellow  substance  produced  by  the  action  of  a  soluble  lead 
salt  upon  a  soluble  chromate,  thus  : 

K2CrO4  +  Pb(NO3)2  =  PbCrO4  +  2  KNO3. 

It  is  used  as  a  yellow  pigment.  Lead  sulphate  (PbSO4)  is 
a  white  substance  sometimes  found  in  nature  and  easily 
prepared  by  precipitation.  Lead  chloride  (PbCl2)  is  like- 
wise a  white  substance  nearly  insoluble  in  cold  water,  but 
readily  soluble  in  boiling  water. 

Thorium  and  cerium.  These  elements  are  found  in  a  few  rare 
minerals,  especially  in  the  monazite  sand  of  the  Carolinas  and  Brazil. 
The  oxides  of  these  elements  are  used  in  the  preparation  of  the  Wels- 
bach  mantles  for  gas  lights,  because  of  the  intense  light  given  out 
when  a  mixture  of  the  oxides  is  heated.  These  mantles  contain  the 
oxides  of  cerium  and  thorium  in  the  ratio  of  about  i%  of  the  former  to 
99%  of  the  latter.  Compounds  of  thorium,  like  those  of  radium,  are 
found  to  possess  radio-activity,  but  in  a  less  degree. 

EXERCISES 

1.  How  could  you  detect  lead  if  present  in  tin  foil? 

2.  Stannous  chloride  reduces  gold  chloride  (AuCl3)  to  gold.    Give 
equation. 

3.  What  are  the  products  of  hydrolysis  when  stannic  chloride  is 
used  as  a  mordant  ? 

4.  How  could  you  detect  arsenic,  antimony,  or  copper  in  lead  ? 


378     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

5.  Why  is  lead  so  extensively  used  for  making  water  pipes  ? 

6.  What  sulphates  other  than  lead  are  insoluble  ? 

7.  Could  lead  nitrate  be  used  in  place  of  barium  chloride  in  te'sting 
for  sulphates  ? 

8.  How  much  lead  peroxide  could  be  obtained  from  i"  kg.  of 
minium  ? 

9.  The  purity  of  white  lead  is  usuall|r  determined  by  observing  the 
volume  of  carbon  dioxide  given  off  when  it  is  treated  with  an  acid. 
What  acid  should  be  used?    On  the  supposition  that  it   has  the 
formula  2PbCO3  •  Pb(OH)2,  how  nearly  pure  was  a  sample  if  I  g. 
gave  30  cc.  of  carbon  dioxide  at  20°  and  750  mm.  ? 

10.  Silicon  belongs  in  the  same  family  with  tin  and  lead.    In  what 
respects  are  these  elements  similar  ? 

11.  What  weight  of  tin  could  be  obtained  by  the  reduction  of  I  ton 
of  cassiterite? 

12.  What  reaction  would  you  expect  to  take  place  when  lead 
peroxide  is  treated  with  hydrochloric  acid  ? 

13.  White  lead  is  often  adulterated  with  barytes.   Suggest  a  method 
for  detecting  it,  if  present,  in  a  given  example  of  white  lead. 


CHAPTER    XXX 
MANGANESE  AND  CHROMIUM 


SYMBOL 

ATOMIC: 
WEIGBH- 

DENSITY 

MELTING 
POINT 

FORMULAS  OF  ACIDS 

Manganese 

Mn 

SS-o  - 

8.6, 

1900° 

.H2MnO4  and   HMnO4 

Chromium 

Cr 

52.1 

7-3 

3000° 

H2CrO4  and   H2Cr2O7 

General.  Manganese  and  chromium,  while  belonging  to 
different  families,  have  so  many  features  in  common  in 
their  chemical  conduct  that  they  may  be  studied  together 
with  advantage.  They  differ  from  most  of  the  elements  so 
far  studied  in  that  they  can  act  either  as  acid-forming  or  base- 
forming  elements.  As  base-forming  elements  each  of  the 
metals  forms  two  series  of  salts.  In  the  one  series,  desig- 
nated by  the  suffix  "  ous,"  the  metal  is  divalent ;  in  the  other 
series,  designated  by  the  suffix  "  ic,"  the  metal  is  trivalent. 
Only  the  manganous  and  the  chromic  salts,  however,  are 
of  importance.  The  acids  in  which  these  elements  play 
the  part  of  a  non-metal  are  unstable,  but  their  salts  are 
usually  stable,  and  some*  of  them  are  important  compounds. 


MANGANESE 

Occurrence.  Manganese  is  found  in  nature  chiefly  as  the 
dioxide  MnO2,  called  pyrolusite.  In  smaller  amounts  it 
occurs  as  the  oxides  Mn2O3  and  Mn3O4,  and  as  the  car-, 
bonate  MnCO3.  Some  iron  ores  also  contain  manganese.. 

379 


380     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Preparation  and  properties.  The  element  is  difficult  to 
prepare  in  pure  condition  and  has  no  commercial  applica- 
tions. It  can  be  prepared,  however,  by  reducing  the  oxide 
with  aluminium  powder  or  by  the  use  of  the  electric 
furnace,  with  carbon  as  the  reducing  agent.  The  metal 
somewhat  resembles  iron  in  appetrance,  but  is  harder,  less 
fusible,  and  more  readily  acted  upon  by  air  and  moisture. 
Acids  readily  dissolve  it,  forming  manganous  salts. 

Oxides  of  manganese.  The  following  oxides  of  manga- 
nese are  known  :  MnO,  Mn2O3,  Mn3O4,  MnO2,  and  Mn2O7. 
Only  one  of  these,  the  dioxide,  needs  special  mention. 

Manganese  dioxide  (pyrolusite)  (MnO2).  This  substance 
is  the  most  abundant  manganese  compound  found  in  nature, 
and  is  the  ore  from  which  all  other  compounds  of  manga- 
nese are  made.  It  is  a  hard,  brittle,  black  substance  which 
is  valuable  as  an  oxidizing  agent.  It  will  be  recalled  that 
it  is  used  in  the  preparation  of  chlorine  and  oxygen,  in 
decolorizing  glass  which  contains  iron,  and  in  the  manufac- 
ture of  ferromanganese. 

Compounds  containing  manganese  as  a  base-forming  ele- 
ment. As  has  been  stated  previously,  manganese  forms 
two  series  of  salts.  The  most  important  of  these  salts,  all 
of  which  belong  to  the  manganous  series,  are  the  following  : 

Manganous  chloride       ....  MnCl2-4H2O. 

Manganous  sulphide      ....  MnS. 

Manganous  sulphate      ....  MnSO4-4.H2O. 

Manganous  carbonate   .     .     .     .  MnCO3. 

Manganous  hydroxide  ....  Mn(OH)2. 

The  chloride  and  sulphate  may  be  prepared  by  heating  the 
dioxide  with  hydrochloric  and  sulphuric  acids  respectively  ; 

MnO2  +  4  HC1  =  MnCl2  +  2  H2O  +  2  Cl, 
MnO2  +  H2SO4  =  MnSO4  +  H2O  +  O. 


MANGANESE  AND  CHROMIUM  381 

The  sulphide,  carbonate,  and  hydroxide,  being  insoluble, 
may  be  prepared  from  a  solution  of  the  chloride  or  sulphate 
by  precipitation  with  the  appropriate  reagents.  Most  of 
the  manganous  salts  are  rose  colored.  They  not  only  have 
formulas  similar  to  the  ferrous  salts,  but  resemble  them  in 
many  of  their  chemical  properties. 

Compounds  containing  manganese  as  an  acid-forming 
element.  Manganese  forms  two  unstable  acids,  namely, 
manganic  acid  and  permanganic  acid.  While  these  acids 
are  of  little  interest,  some  of  their  salts,  especially  the 
permanganates,  are  important  compounds. 

Manganic  acid  and  manganates.  When  manganese  diox- 
ide is  fused  with  an  alkali  and  an  oxidizing  agent  a  green 
compound  is  formed.  The  equation,  when  caustic  potash 
is  used,  is  as  follows  : 

MnO2  +  2  KOH  +  O  =  K2MnO4  +  H2O. 

The  green  compound  (K2MnO4)  is  called  potassium 
manganate,  and  is  a  salt  of  the  unstable  manganic  acid 
(H2MnO4).  The  manganates  are  all  very  unstable. 

Permanganic  acid  and  the  permanganates.  When  carbon 
dioxide  is  passed  through  a  solution  of  a  manganate  a  part 
of  the  manganese  is  changed  into  manganese  dioxide,  while 
the  remainder  forms  a  salt  of  the  unstable  acid  HMnO4, 
called  permanganic  acid.  The  equation*  is 

3  K2MnO4  +  2  CO2  =  MnO2  +  2  KMnO4  +  2  K2CO3. 

Potassium  permanganate  (KMnO4)  crystallizes  in  purple- 
black  needles  and  is  very  soluble  in  water,  forming  an 
intensely  purple  solution.  All  other  permanganates,  as 
well  as  permanganic  acid  itself,  give  solutions  of  the  same 
color. 


382     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Oxidizing  properties  of  the  permanganates.  The  perman- 
ganates are  remarkable  for  their  strong  oxidizing  properties. 
When  used  as  an  oxidizing  agent  the  permanganate  is 
itself  reduced,  the  exact  character  of  the  products  formed 
from  it  depending  upon  whether  the  oxidation  takes  place 
(i)  in  an  alkaline  or  neutral  solution,  or  (2)  in  an  acid 
solution. 

1.  Oxidation  in  alkaline  or  neutral  solution.    When  the 
solution  is  either  alkaline  or  neutral  the  potassium  and  the 
manganese  of  the  permanganate  are  both  converted  into 
hydroxides,  as  shown  in  the  equation 

2  KMnO4  +  5  H2O  =  2  Mn(OH)4  +  2  KOH  +  30. 

2.  Oxidation  in  acid  solution.    When  free  acid  such  as 
sulphuric  is  present,  the  potassium  and  the  manganese  are 
both  changed  into  salts  of  the  acid  : 

2  KMnO4  +  3  H2SO4  -  K2SO4  +  2  MnSO4  +  3  H2O  +  5  O. 

Under  ordinary  conditions,  however,  neither  one  of  these 
reactions  takes  place  except  in  the  presence  of  a  third  sub- 
stance which  is  capable  of  oxidation.  The  oxygen  is  not 
given  off  in  the  free  state,  as  the  equations  show,  but  is 
used  up  in  effecting  oxidation. 

Potassium  permanganate  is  particularly  valuable  as  an 
oxidizing  agent  not  only  because  it  acts  readily  either  in 
acid  or  in  alkaline  solution,  but  also  because  the  reaction 
takes  place  so  easily  that  often  it  is  not  even  necessary 
to  heat  the  solution  to  secure  action.  The  substance 
finds  many  uses  in  the  laboratory,  especially  in  analytical 
work.  It  is  also  used  as  an  antiseptic  as  well  as  a  disin- 
fectant. 


MANGANESE  AND  CHROMIUM  383 

CHROMIUM 

Occurrence.  The  ore  from  which  all  chromium  com- 
pounds are  made  is  chromite,  or  chrome  iron  ore  (FeCr2O4). 
This  is  found  most  abundantly  in  New  Caledonia  and 
Turkey.  The  element  also  occurs  in  small  quantities  in 
many  other  minerals,  especially  in  crocoisite  (PbCrO4),  in 
which  mineral  it  was  first  discovered. 

Preparation.  Chromium,  like  manganese,  is  very  hard  to 
reduce  from  its  ores,  owing  to  its  great  affinity  for  oxygen. 
It  can,  however,  be  made  by  the  same  methods  which 
have  proved  successful  with  manganese.  Considerable 
quantities  of  an  alloy  of  chromium  with  iron,  called  ferro- 
chromium,  are  now  produced  for  the  steel  industry. 

Properties.  Chromium  is  a  very  hard  metal  of  about  the 
same  density  as  iron.  It  is  one  of  the  most  infusible  of  the 
metals,  requiring  a  temperature  little  short  of  3000°  for 
fusion.  At  ordinary  temperatures  air  has  little  action  on 
it ;  at  higher  temperatures,  however,  it  burns  brilliantly. 
Nitric  acid  has  no  action  on  it,  but  hydrochloric  and  dilute 
sulphuric  acids  dissolve  it,  liberating  hydrogen. 

Compounds  containing  chromium  as  a  base-forming 
element.  While  chromium  forms  two  series  of  salts, 
chromous  salts  are  difficult  to  prepare  and  are  of  little 
importance.  The  most  important  of  the  chromic  series 
are  the  following  : 

Chromic  hydroxide  ......  Cr(OH)3. 

Chromic  chloride CrCl3  •  6H2O. 

Chromic  sulphate Cr2(SO4)3. 

Chrome  alums 

Chromic  hydroxide  (Cr(OH)3).  This  substance,  being 
insoluble,  can  be  obtained  by  precipitating  a  solution  of 
the  chloride  or  sulphate  with  a  soluble  hydroxide.  It  is  a 


384     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

greenish  substance  which,  like  aluminium  hydroxide,  dis- 
solves in  alkalis,  forming  soluble  salts. 

Dehydration  of  chromium  hydroxide.  When  heated  gently  chromic 
hydroxide  loses  a  part  of  its  oxygen  and  hydrogen,  forming  the  sub- 
stance CrO  •  OH,  which,  like  the  corresponding  aluminium  compound, 
has  more  pronounced  acid  properties  than  the  hydroxide.  It  forms  a 
series  of  salts  very  similar  to  the  spinels  ;  chromite  is  the  ferrous 
salt  of  this  acid,  having  the  formula  Fe(CrO2)2.  When  heated  to  a 
higher  temperature  chromic  hydroxide  is  completely  dehydrated,  form- 
ing the  trioxide  Cr2O3.  This  resembles  the  corresponding  oxides  of 
aluminium  and  iron  in  many  respects.  It  is  a  bright  green  powder, 
and  when  ignited  strongly  becomes  almost  insoluble  in  acids,  as  is 
also  the  case  with  aluminium  oxide. 

Chromic  sulphate  (Cr2(SO4)3).  This  compound  is  a  vio- 
let-colored solid  which  dissolves  in  water,  forming  a  solu- 
tion of  the  same  color.  This  solution,  however,  turns 
green  on  heating,  owing  to  the  formation  of  basic 
salts.  Chromic  sulphate,  like  ferric  and  aluminium  sul- 
phates, unites  with  the  sulphates  of  the  alkali  metals  to 
form  alums,  of  which  the  best  known  are  potassium 
chrome  alum  (KCr(SO4)2  -  12  H2O)  and  ammonium  chrome 
alum  (NH4Cr(SO4)2-  12  H2O). 

These  form  beautiful  dark  purple  crystals  and  have  some 
practical  uses  in  the  tanning  industry  and  in  photography. 
A  number  of  the  salts  of  chromium  are  also  used  in  the 
dyeing  industry,  for  they  hydrolyze  like  aluminium  salts 
and  the  hydroxide  forms  a  good  mordant. 

Hydrolysis  of  chromium  salts.  When  ammonium  sulphide  is  added 
to  a  solution  of  a  chromium  salt,  such  as  the  sulphate,  chromium 
hydroxide  precipitates  instead  of  the  sulphide.  This  is  due  to  the 
fact  that  chromic  sulphide,  like  aluminium  sulphide,  hydrolyzes  in  the 
presence  of  water,  forming  chromic  hydroxide  and  hydrosulphuric 
acid.  Similarly,  a  soluble  carbonate  precipitates  a  basic  carbonate 
of  chromium. 


MANGANESE  AND  CHROMIUM  385 

Compounds  containing  chromium  as  an  acid-forming  ele- 
ment. Like  manganese,  chromium  forms  two  unstable 
acids,  namely,  chromic  acid  and  dichromic  acid.  Their 
salts,  the  chromates  and  dichromates,  are  important 
compounds. 

Chromates.  When  a  chromium  compound  is  fused  with 
an  alkali  and  an  oxidizing  agent  a  chromate  is  produced. 
When  potassium  hydroxide  is  used  as  the  alkali  the  equa- 
tion is 

2  Cr(OH)3  +  4  KOH  +  3  O  -  2  K2CrO4  +  5  H2O. 

This  reaction  recalls  the  formation  of  a  manganate  under 
similar  conditions. 

Properties  of  chromates.  The  chromates  are  salts  of  the 
unstable  chromic  acid  (H2CrO4),  and  as  a  rule  are  yellow 
in  color.  Lead  chromate  (PbCrO4)  is  the  well-known  pig- 
ment chrome  yellow.  Most  of  the  chromates  are  insoluble 
and  can  therefore  be  prepared  by  precipitation.  Thus,  when 
a  solution  of  potassium  chromate  is  added  to  solutions  of 
lead  nitrate  and  barium  nitrate  respectively,  the  reactions 
expressed  by  the  following  equations  occur  : 

Pb(NO3)2  +  K2CrO4  -  PbCrO4  +  2  KNO3, 
Ba(NO3)2  +  K2CrO4  -  BaCrO4  +  2  KNO3. 

The  chromates  of  lead  and  barium  separate  as  yellow  pre- 
cipitates. The  presence  of  either  of  these  two  metals 
can  be  detected  by  taking  advantage  of  these  reactions. 
Dichromates.  When  potassium  chromate  is  treated  with 
an  acid  the  potassium  salt  of  the  unstable  dichromic  acid 
(H2Cr2O7)  is  formed  : 

2  K2CrO4  +  H2SO4  -  K2Cr2O7  +  K2SO4  +  H2O. 


386     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

The  relation  between  the  chromates  and  dichromates  is 
the  same  as  that  between  the  phosphates  and  the  pyro- 
phosphates.  Potassium  dichromate  might  therefore  be 
called  potassium  pyrochromate. 

Potassium  dichromate  (K2Cr2O7).  This  is  the  best  known 
dichromate,  and  is  the  most  familiar  chromium  compound. 
It  forms  large  crystals  of  a  brilliant  red  color,  and  is  rather 
sparingly  soluble  in  water.  When  treated  with  potassium 
hydroxide  it  is  converted  into  the  chromate 

K2Cr2O7  +  2  KOH=  2  K2CrO4  +  H2O. 

When  added  to  a  solution  of  lead  or  barium  salt  the  corre- 
sponding chromates  (not  dichromates)  are  precipitated. 
With  barium  nitrate  the  equation  is 

2  Ba(NO3)2  +  K2Cr2O7  +  H2O  =  2  BaCrO4  + 
2KNO3+2HNO3. 

Potassium  dichromate  finds  use  in  many  industries  as  an 
oxidizing  agent,  especially  in  the  preparation  of  organic 
substances,  such  as  the  dye  alizarin,  and  in  the  construc- 
tion of  several  varieties  of  electric  batteries. 

Sodium  chromates.  The  reason  why  the  potassium  salt  rather 
than  the  sodium  compound  is  used  is  that  sodium  chromate  and 
dichromate  are  so  soluble  that  it  is  hard  to  prepare  them  pure. 
This  difficulty  is  being  overcome  now,  and  the  sodium  compounds 
are  replacing  the  corresponding  potassium  salts.  This  is  of  advantage, 
since  a  sodium  salt  is  cheaper  than  a  potassium  salt,  so  far  as  raw 
materials  go. 

Oxidizing  action  of  chromates  and  dichromates.  When  a 
dilute  solution  of  a  chromate  or  dichromate  is  acidified 
with  an  acid,  such  as  sulphuric  acid,  no  reaction  apparently 
takes  place.  However,  if  there  is  present  a  third  substance 
capable  of  oxidation,  the  chromium  compound  gives  up  a 


MANGANESE  AND  CHROMIUM  387 

portion  of  its  oxygen  to  this  substance.  Since  the  chromate 
changes  into  a  dichromate  in  the  presence -of  an  acid,  it 
will  be  sufficient  to  study  the  action  of  the  dichromates 
alone.  The  reaction  takes  place  in  two  steps.  Thus,  when 
a  solution  of  ferrous  sulphate  is  added  to  a  solution  of 
potassium  dichromate  acidified  with  sulphuric  acid,  the 
reaction  is  expressed  by  the  following  equations  : 

(1)  K2Cr207  +  4  H2S04  =  K2SO4  +  Cr2(SO4)3  + 

4  H20  +  3  O, 

(2)  6  FeS04  +  3  H2SO4  +30  =  3  Fe2(SO4)3  +  3  H2O. 

The  dichromate  decomposes  in  very  much  the  same  way 
as  a  permanganate  does,  the  potassium  and  chromium 
being  both  changed  into  salts  in  which  they  play  the  part 
of  metals,  while  part  of  the  oxygen  of  the  dichromate  is 
liberated. 

By  combining  equations  (i)  and  (2),  the  following  is 
obtained : 

K2Cr207  +  7  H2S04  +  6  FeSO4  =  K2SO4  +  Cr2(SO4)3  + 
3  Fe2(S04)3  +  7  H20. 

This  reaction  is  often  employed  in  the  estimation  of  iron 
in  iron  ores. 

Potassium  chrome  alum.  It  will  be  noticed  that  the  oxidizing 
action  of  potassium  dichromate  leaves  potassium  sulphate  and  chro- 
mium sulphate  as  the  products  of  the  reaction'.  On  evaporating  the 
solution  these  substances  crystallize  out  as  potassium  chrome  alum, 
which  substance  is  produced  as  a  by-product  in  the  industries  using 
potassium  dichromate  for  oxidizing  purposes. 

Chromic  anhydride  (CrO3).  When  concentrated  sulphuric 
acid  is  added  to  a  strong  solution  of  potassium  dichromate, 
and  the  liquid  allowed  to  stand,  deep  red  needle-shaped 
crystals  appear  which  have  the  formula  CrO3.  This  oxide 


388     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

of  chromium  is  called  chromic  anhydride,  since  it  combines 
readily  with  water  to  form  chromic  acid : 

CrO3  +  H2O  -  H2CrO4. 

It  is  therefore  analogous  to  sulphur  trioxide  which  forms 
sulphuric  acid  in  a  similar  way: 

SO3  +  H2O  -  H2SO4. 

Chromic  anhydride  is  a  very  strong  oxidizing  agent,  giving 
up  oxygen  and  forming  chromic  oxide  : 

2  CrO3  -  Cr2O3  +  3  O. 

Rare  elements  of  the  family.  Molybdenum,  tungsten,  and  uranium 
are  three  rather  rare  elements  belonging  in  the  same  family  with 
chromium,  and  form  many  compounds  which  are  similar  in  formulas 
to  the  corresponding  compounds  of  chromium.  They  can  play  the 
part  of  metals  and  also  form  acids  resembling  chromic  acid  in 
formula.  Thus  we  have  molybdic  acid  (H2MoO4),  the  ammonium 
salt  of  which  is  (NH4)2MoO4.  This  salt  has  the  property  of  com- 
bining with  phosphoric  acid  to  form  a  very  complex  substance  which 
is  insoluble  in  nitric  acid.  On  this  account  molybdic  acid  is  often 
used  in  the  estimation  of  the  phosphoric  acid  present  in  a  substance. 
Like  chromium,  the  metals  are  difficult  to  prepare  in  pure  condition. 
Alloys  with  iron  can  be  prepared  by  reducing  the  mixed  oxides  with 
carbon  in  an  electric  furnace ;  these  alloys  are  used  to  some  extent 
in  preparing  special  kinds  of  steel. 


EXERCISES 

1.  How  does  pyrolusite  effect  the  decolorizing  of  glass  containing 
iron  ? 

2.  Write  the  equations  for  the  preparation  of  manganous  chloride, 
carbonate,  and  hydroxide. 

3.  Write  the  equations  representing  the  reactions  which  take  place 
when  ferrous  sulphate  is  oxidized  to  ferric  sulphate  by  potassium  per- 
manganate in  the  presence  of  sulphuric  acid. 


MANGANESE  AND  CHROMIUM  389 

4.  In  the  presence  of  sulphuric  acid,  oxalic  acid  is  oxidized  by 
potassium  permanganate  according  to  the  equation 

C2H2O4  +  O  =  2  CO2  +  H2O. 

Write  the  complete  equation. 

5.  log.  of  iron  were  dissolved  in  sulphuric  acid  and  oxidized  to 
ferric  sulphate  by  potassium  permanganate.    What  weight  of  the 
permanganate  was  required  ? 

6.  What  weight  of  ferrochromium  containing  40%  chromium  must 
be  added  to  a  ton  of  steel  to  produce  an  alloy  containing  i%  of 
chromium  ? 

7.  Write  the  equation  representing  the  action  of  ammonium  sul- 
phide upon  chromium  sulphate. 

8.  Potassium  chromate  oxidizes  hydrochloric  acid,  forming  chlorine. 
Write  the  complete  equation. 

9.  Give  the  action  of  sulphuric  acid  on  potassium  dichromate  (#) 
in  the  presence  of  a  large  amount  of  water ;  (£)  in  the  presence  of  a 
small  amount  of  water. 


CHAPTER   XXXI 
GOLD  AND  THE  PLATINUM  FAMILY 


SYMBOL 

ATOMIC 
WEIGHT 

DENSITY 

HIGHEST 
OXIDE 

HIGHEST 
CHLORIDE 

MELTING 
POINT 

Ruthenium   . 

Ru 

100.9 

12.26 

RuO4 

RuCl4 

Electric  arc 

Rhodium  .     . 

Rh 

IO2.2 

12.  1 

RhO2 

RhCl3 

Electric  arc 

Palladium     . 

Pd 

105.2 

II.8 

PdO2 

PdCl4 

1500° 

Iridium     .     . 

Ir 

191-S 

22.42 

IrO2 

IrCl4 

1950° 

Osmium   .     . 

Os 

189.6 

22.47 

OsO4 

OsCl4 

Electric  arc 

Platinum  .     . 

Pt 

193-3 

21.50 

Pt02 

PtCl4 

1779° 

Gold    .     .     . 

Au 

195-7 

19.30 

Au2O3 

AuCl3 

1064° 

The  family.  Following  iron,  nickel,  and  cobalt  in  the 
eighth  column  of  the  periodic  table  are  two  groups  of 
three  elements  each.  The  metals  of  the  first  of  these 
groups — ruthenium,  rhodium,  and  palladium — have  atomic 
weights  near  100  and  densities  near  12.  The  metals  of 
the  other  group  —  iridium,  osmium,  and  platinum  —  have 
atomic  weights  near  200  and  densities  near  21.  These 
six  rare  elements  have  very  similar  physical  properties 
and  resemble  each  other  chemically  not  only  in  the  type  of 
compounds  which  they  form  but  also  in  the  great  variety 
of  them.  They  occur  closely  associated  in  nature,  usually 
as  alloys  of  platinum  in  the  form  of  irregular  metallic 
grains  in  sand  and  gravel.  Platinum  is  by  far  the  most 
abundant  of  the  six. 

Although  the  periodic  classification  assigns  gold  to  the 
silver-copper  group,  its  physical  as  well  as  many  of  its 

390 


GOLD  AND  THE  PLATINUM  FAMILY  391 

chemical  properties  much  more  closely  resemble  those  of 
the  platinum  metals,  and  it  can  be  conveniently  considered 
along  with  them.  The  four  elements  gold,  platinum,  osmi- 
um, and  iridium  are  the  heaviest  substances  known,  being 
about  twice  as  heavy  as  lead. 

PLATINUM 

Occurrence.  About  90%  of  the  platinum  of  commerce 
comes  from  Russia,  small  amounts  being  produced  in 
California,  Brazil,  and  Australia. 

Preparation.  Native  platinum  is  usually  alloyed  with 
gold  and  the  platinum  metals.  To  separate  the  platinum 
the  alloy  is  dissolved  in  aqua  regia,  which  converts  the 
platinum  into  chloroplatinic  acid  (H2PtCl6).  Ammonium 
chloride  is  then  added,  which  precipitates  the  platinum  as 
insoluble  ammonium  chloroplatinate  : 

H2PtCl6  +  2.  NH4C1  =  (NH4)2PtCl6  +  2  HC1. 

Some  iridium  is  also  precipitated  as  a  similar  compound. 
On  ignition  the  double  chloride  is  decomposed,  leaving  the 
platinum  as  a  spongy  metallic  mass,  which  is  melted  in  an 
electric  furnace  and  rolled  or  hammered  into  the  desired 
shape. 

Physical  properties.  Platinum  is  a  grayish-white  metal 
of  high  luster,  and  is  very  malleable  and  ductile.  It  melts 
in  the  oxyhydrogen  blowpipe  and  in  the  electric  furnace ; 
it  is  harder  than  gold  and  is  a  good  conductor  of  electricity. 
In  finely  divided  form  it  has  the  ability  to  absorb  or  occlude 
gases,  especially  oxygen  and  hydrogen.  These  gases,  when 
occluded,  are  in  a  very  active  condition  resembling  the 
nascent  state,  and  can  combine  with  each  other  at  ordinary 


392     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

temperatures.    A  jet  of  hydrogen  or  coal  gas  directed  upon 
spongy  platinum  is  at  once  ignited. 

Platinum  as  a  catalytic  agent.  Platinum  is  remarkable  for  its  prop- 
erty of  acting  as  a  catalytic  agent  in  a  large  number  of  chemical 
reactions,  and  mention  has  been  made  of  this  use  of  the  metal  in  con- 
nection with  the  manufacture  of  sulphuric  acid.  When  desired  for 
this  purpose  some  porous  or  fibrous  substance,  such  as  asbestos,  is 
soaked  in  a  solution  of  platinic  chloride  and  then  ignited.  The  plat- 
inum compound  is  decomposed  and  the  platinum  deposited  in  very 
finely  divided  form.  Asbestos  prepared  in  this  way  is  called  plat- 
inized asbestos.  The  catalytic  action  seems  to  be  in  part  connected 
with  the  property  of  absorbing  gases  and  rendering  them  nascent. 
Some  other  metals  possess  this  same  power,  notably  palladium, 
which  is  remarkable  for  its  ability  to  absorb  hydrogen. 

Chemical  properties.  Platinum  is  a  very  inactive  element 
chemically,  and  is  not  attacked  by  any  of  the  common 
acids.  Aqua  regia  slowly  dissolves  it,  forming  platinic 
chloride  (PtCl4),  which  in  turn  unites  with  the  hydro- 
chloric acid  present  in  the  aqua  regia,  forming  the  com- 
pound chloroplatinic  acid  (H2PtCl6).  Platinum  is  attacked 
by  fused  alkalis.  It  combines  at  higher  temperatures  with 
carbon  and  phosphorus  and  alloys  with  many  metals.  It 
is  readily  attacked  by  chlorine  but  not  by  oxidizing 
agents. 

Applications.  Platinum  is  very  valuable  as  a  material  for 
the  manufacture  of  chemical  utensils  which  are  required 
to  stand  a  high  temperature  or  the  action  of  strong  re- 
agents. Platinum  crucibles,  dishes,  forceps,  electrodes,  and 
similar  articles  are  indispensable  in  the  chemical  laboratory. 
In  the  industries  it  is  used  for  such  purposes  as  the  man- 
ufacture of  pans  for  evaporating  sulphuric  acid,  wires  for 
sealing  through  incandescent  light  bulbs,  and  for  making 
a  great  variety  of  instruments.  Unfortunately  the  supply 


GOLD  AND  THE  PLATINUM  FAMILY  393 

of  the  metal  is  very  limited,  and  the  cost  is  steadily 
advancing,  so  that  it  is  now  more  valuable  than  gold. 

Compounds.  Platinum  forms  two  series  of  salts  of  which 
platinous  chloride  (PtCl2)  and  platinic  chloride  (PtCl4) 
are  examples.  Platinates  are  also  known.  While  a  great 
variety  of  compounds  of  platinum  have  been  made,  the 
substance  is  chiefly  employed  in  the  metallic  state. 

Platinic  chloride  (PtCl4).  Platinic  chloride  is  an  orange- 
colored,  soluble  compound  made  by  heating  chloroplatinic 
acid  in  a  current  of  chlorine.  If  hydrochloric  acid  is  added 
to  a  solution  of  the  substance,  the  two  combine,  forming 
chloroplatinic  acid  (H2PtCl6) : 

2  HC1  +  PtCl4  -  H2PtCl6. 

The  potassium  and  ammonium  salts  of  this  acid  are  nearly 
insoluble  in  water  and  alcohol.  The  acid  is  therefore  used 
as  a  reagent  to  precipitate  potassium  in  analytical  work. 
With  potassium  chloride  the  equation  is 

2  KC1  +  H2PtCl6  =  K2PtCl6  +  2  HC1. 

Other  metals  of  the  family.  The  other  members  of  the  family  have 
few  applications.  Iridium  is  used  in  the  form  of  a  platinum  alloy, 
since  the  alloy  is  much  harder  than  pure  platinum  and  is  even  less 
fusible.  This  alloy  is  sometimes  used  to  point  gold  pens.  Osmium 
tetroxide  (OsO4)  is  a  very  volatile  liquid  and  is  used  under  the  name 
of  osmic  acid  as  a  stain  for  sections  in  microscopy. 

GOLD 

Occurrence.  Gold  has  been  found  in  many  localities,  the 
most  famous  being  South  Africa,  Australia,  Russia,  and 
the  United  States.  In  this  country  it  is  found  in  Alaska 
and  in  nearly  half  of  the  states  of  the  union,  notably 


394     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

in  California,  Colorado,  and  Nevada.  It  is  usually  found 
in  the  native  condition,  frequently  alloyed  with  silver ;  in 
combination  it  is  sometimes  found  as  telluride  (AuTe2), 
and  in  a  few  other  compounds. 

Mining.  Native  gold  occurs  in  the  form  of  small  grains 
or  larger  nuggets  in  the  sands  of  old  rivers,  or  imbedded 
in  quartz  veins  in  rocks.  In  the  first  case  it  is  obtained  in 
crude  form  by  placer  mining.  The  sand  containing  the 
gold  is  shaken  or  stirred  in  troughs  of  running  waters 
called  sluices.  This  sweeps  away  the  sand  but  allows  the 
heavier  gold  to  sink  to  the  bottom  of  the  sluice.  Some- 
times the  sand  containing  the  gold  is  washed  away  from 
its  natural  location  into  the  sluices  by  powerful  streams 
of  water  delivered  under  pressure  from  pipes.  This  is 
called  hydraulic  mining.  In  vein  mining  the  gold-bearing 
quartz  is  mined  from  the  veins,  stamped  into  fine  powder 
in  stamping  mills,  and  the  gold  extracted  by  one  of  the 
processes  to  be  described. 

Extraction,  i.  Amalgamation  process.  In  the  amalga- 
mation process  the  powder  containing  the  gold  is  washed 
over  a  series  of  copper  plates  whose  surfaces  have  been 
amalgamated  with  mercury.  The  gold  sticks  to  the  mer- 
cury or  alloys  with  it,  and  after  a  time  the  gold  and 
mercury  are  scraped  off  and  the  mixture  is  distilled. 
The  mercury  distills  off  and  the  gold  is  left  in  the  retort 
ready  for  refining. 

2.  Chlorination  process.  When  gold  occurs  along  with 
metallic  sulphides  it  is  often  extracted  by  chlorination. 
The  ore  is  first  roasted,  and  is  then  moistened  and  treated 
with  chlorine.  This  dissolves  the  gold  but  not  the  metallic 
oxides  : 

Au  +  3  Cl  =  AuCl3. 


GOLD  AND  THE  PLATINUM  FAMILY  395 

The  gold  chloride,  being  soluble,  is  extracted  from  the 
mixture  with  water,  and  the  gold  is  precipitated  from  the 
solution,. usually  by  adding  ferrous  sulphate  : 

AuCl3  +  3  FeSO4  =  Au  +  FeCl3  +  Fe2(SO4)3. 

3.  Cyanide  process.  This  process  depends  upon  the  fact 
that  gold  is  soluble  in  a  solution  of  potassium  cyanide  in 
the  presence  of  the  oxygen  of  the  air.  The  powder  from 
the  stamping  mills  is  treated  with  a  very  dilute  potassium 
cyanide  solution  which  extracts  the  gold  : 

2  Au  +  4  KCN  +  H2O  +  O  =  2  KOH  +  2  KAu(CN)2. 

From  this  solution  the  gold  can  be  obtained  by  electrolysis 
or  by  precipitation  with  metallic  zinc  : 

2  KAu(CN)2  +  Zn  =  K2Zn(CN)4  +  2  Au. 

Refining  of  gold.  Gold  is  refined  by  three  general 
methods  : 

1.  Electrolysis.    When  gold  is  dissolved  in  a  solution  of 
potassium  cyanide,  and  the  solution  electrolyzed,  the  gold 
is  deposited  in  very  pure  condition  on  the  anode. 

2.  Cupellation.    When  the   gold   is   alloyed  with  easily 
oxidizable  metals,  such  as  copper  or  lead,  it  may  be  refined 
by  cupellation.    The  alloy  is  fused  with  an  oxidizing  flame 
on  a  shallow  hearth  made  of  bone  ash,  which  substance  has 
the  property  of  absorbing  metallic  oxides  but  not  the  gold. 
Any  silver  which  may  be  present  remains  alloyed  with  the 
gold. 

3.  Parting  with  sulphuric  acid.    Gold  may  be  separated 
from  silver,  as  well  as  from  many  other  metals,  by  heating 
the  alloy  with  concentrated  sulphuric  acid.    This  dissolves 
the  silver,  while  the  gold  is  not  attacked. 


396     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Physical  properties.  Gold  is  a  very  heavy  bright  yellow 
metal,  exceedingly  malleable  and  ductile,  and  a  good  con- 
ductor of  electricity.  It  is  quite  soft  and  is  usually  alloyed 
with  copper  or  silver  to  give  it  the  hardness  required  for 
most  practical  uses.  The  degree  of  fineness  is  expressed 
in  terms  of  carats,  pure  gold  being  twenty-four  carats; 
the  gold  used  for  jewelry  is  usually  eighteen  carats, 
eighteen  parts  being  gold  and  six  parts  copper  or  silver. 
Gold  coinage  is  90%  gold  and  10%  copper. 

Chemical  properties.  Gold  is  not  attacked  by  any  one  of 
the  common  acids ;  aqua  regia  easily  dissolves  it,  forming 
gold  chloride  (AuCl3),  which  in  turn  combines  with  hydro- 
chloric acid  to  form  chlorauric  acid  (HAuCl4).  Fused 
alkalis  also  attack  it.  Most  oxidizing  agents  are  without 

action  upon  it,  and  in  general  it  is  not  an  active  element. 

/ 

Compounds.  The  compounds  of  gold,  though  numerous  and  varied 
in  character,  are  of  comparatively  little  importance  and  need  not  be 
described  in  detail.  The  element  forms  two  series  of  salts  in  which 
it  acts  as  a  metal:  in  the  aurous  series  the  gold  is  univalent,  the 
chloride  having  the  formula  AuCl;  in  the  auric  series  it  is  trivalent, 
auric  chloride  having  the  formula  AuCl3.  Gold  also  acts  as  an 
acid-forming  element,  forming  such  compounds  as  potassium  aurate 
(KAuO2).  Its  compounds  are  very  easily  decomposed,  however, 
metallic  gold  separating  from  them. 

EXERCISES 

1.  From  the  method  of  preparation  of  platinum,  what  metal  is  likely 
to  be  alloyed  with  it  ? 

2.  The  "  platinum  chloride  "  of  the  laboratory  is  made  by  dissolving 
platinum  in  aqua  regia.    What  is  the  compound? 

3.  How  would  you  expect  potassium  aurate  and  platinate  to  be 
formed  ?    What  precautions  would  this  suggest  in  the  use  of  platinum 
vessels  ? 

4.  Why  must  gold  ores  be  roasted  in  the  chlorination  process  ? 


CHAPTER   XXXII 
SOME  SIMPLE  ORGANIC  COMPOUNDS 

Division  of  chemistry  into  organic  and  inorganic.  Chem- 
istry is  usually  divided  into  two  great  divisions,  —  organic 
and  inorganic.  The  original  significance  of  these  terms 
was  entirely  different  from  the  meaning  which  they  have  at 
the  present  time. 

i.  Original  significance .  The  division  into  organic  and 
inorganic  was  originally  made  because  it  was  believed  that 
those  substances  which  constitute  the  essential  parts  of 
living  organisms  were  built  up  under  the  influence  of  the 
life  force  of  the  organism.  Such  substances,  therefore, 
should  be  regarded  as  different  from  those  compounds  pre- 
pared in  the  laboratory  or  formed  from  the  inorganic  or 
mineral  constituents  of  the  earth.  In  accordance  with 
this  view  organic  chemistry  included  those  substances 
formed  by  living  organisms.  Inorganic  chemistry,  on  the 
other  hand,  included  all  substances  formed  from  the 
mineral  portions  of  the  earth. 

In  1828  the  German  chemist  Wohler  prepared  urea, 
a  typical  organic  compound,  from  inorganic  materials. 
The  synthesis  of  other  so-called  organic  compounds  fol- 
lowed, and  at  present  it  is  known  that  the  same  chemical 
laws  apply  to  all  substances  whether  formed  in  the  living 
organism  or  prepared  in  the  laboratory  from  inorganic  con- 
stituents. The  terms  "  organic  "  and  "  inorganic"  have 
therefore  lost  their  original  significance. 

397 


398     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

2.  Present  significance.  The  great  majority  of  the  com- 
pounds found  in  living  organisms  contain  carbon,  and  the 
term  "  organic  chemistry,"  as  used  at  present,  includes 
not  only  these  compounds  but  all  compounds  of  carbon. 
Organic  chemistry  has  become,  therefore,  the  chemistry 
of  the  compounds  of  carbon,  all  other  substances  being 
treated  under  the  head  of  inorganic  chemistry.  This  sepa- 
ration of  the  compounds  of  carbon  into  a  group  by  them- 
selves is  made  almost  necessary  by  their  great  number, 
over  one  hundred  thousand  having  been  recorded.  For 
convenience  some  of  the  simpler  carbon  compounds,  such 
as  the  oxides  'and  the  carbonates,  are  usually  discussed  in 
inorganic  chemistry. 

The  grouping  of  compounds  in  classes.  The  study  of 
organic  chemistry  is  much  simplified  by  the  fact  that  the 
large  number  of  bodies  included  in  this  field  may  be  grouped 
in  classes  of  similar  compounds.  It  thus  becomes  possible 
to  study  the  properties  of  each  class  as  a  whole,  in  much 
the  same  way  as  we  study  a  group  of  elements.  The  most 
important  of  these  classes  are  the  hydrocarbons,  the  alco- 
hols, the  aldehydes,  the  acids,  the  ethereal  salts,  the  ethers, 
the  ke tones,  the  organic  bases,  and  the  carbohydrates.  A  few 
members  of  each  of  these  classes  will  now  be  discussed 
briefly. 

THE  HYDROCARBONS 

Carbon  and  hydrogen  combine  to  form  a  large  number 
of  compounds.  These  compounds  are  known  collectively  as 
the  hydrocarbons.  They  may  be  divided  into  a  number  of 
groups  or  series,  each  being  named  from  its  first  member. 
Some  of  the  groups  are  as  follows : 


SOME  SIMPLE  ORGANIC  COMPOUNDS         399 

METHANE  SERIES  ETHYLENE  SERIES 

CH4  —  methane  C2H4  —  ethylene 

C2H6  —  ethane  C3H6  —  propylene 

C3H8  —  propane  C4H8  —  butylene 
C4H10  —  butane 

C5H12  -  pentane  BENZENE  SERIES 

C6H14  —  hexane  C6H6  —  benzene 

C7H16  —  heptane  C7H8  —  toluene 

C8H18  -  octane  C8H10—  xylene 

ACETYLENE  SERIES 
C2H2  —  acetylene 
C3H4  -  allylene 

Only  the  lower  members  (that  is,  those  which  contain 
a  small  number  of  carbon  atoms)  of  the  above  groups  are 
given.  The  methane  series  is  the  most  extensive,  all  of 
the  compounds  up  to  C24H50  being  known. 

It  will  be  noticed  that  the  successive  members  of  each 
of  the  above  series  differ  by  the  group  of  atoms  (CH2). 
Such  a  series  is  called  an  homologous  series.  In  general,  it 
may  be  stated  that  the  members  of  an  homologous  series 
show  a  regular  gradation  in  most  physical  properties  and 
are  similar  in  chemical  properties.  Thus  in  the  methane 
group  the  first  four  members  are  gases  at  ordinary 
temperatures  ;  those  containing  from  five  to  sixteen  carbon 
atoms  are  liquids,  the  boiling  points  of  which  increase  with 
the  number  of  carbon  atoms  present..  Those  containing 
more  than  sixteen  carbon  atoms  are  solids. 

Sources  of  the  hydrocarbons.  There  are  two  chief  sources 
of  the  hydrocarbons,  namely,  (i)  crude  petroleum  and  (2) 
coal  tar. 

i.  Crude  petroleum.  This  is  a  liquid  pumped  from 
wells  driven  into  the  earth  in  certain  localities.  Pennsyl- 
vania, Ohio,  Kansas,  California,  and  Texas  are  the  chief 


400     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

oil-producing  regions  in  the  United  States.  The  crude 
petroleum  consists  largely  of  liquid  hydrocarbons  in  which 
are  dissolved  both  gaseous  and  solid  hydrocarbons.  Before 
being  used  it  must  be  refined.  In  this  process  the  petroleum 
is  run  into  large  iron  stills  and  subjected  to  fractional  dis- 
tillation. The  various  hydrocarbons  distill  over  in  the  gen- 
eral order  of  their  boiling  points.  The  distillates  which 
collect  between  certain  limits  of  temperature  are  kept 
separate  and  serve  for  different  uses  ;  they  are  further 
purified,  generally  by  washing  with  sulphuric  acid,  then 
with  an  alkali,  and  finally  with  water.  Among  the  prod- 
ucts obtained  from  crude  petroleum  in  this  way  are  the 
naphthas,  including  benzine  and  gasoline,  kerosene  or  coal 
oil,  lubricating  oils,  vaseline,  and  paraffin.  None  of  these 
products  are  definite  chemical  compounds,  but  each  con- 
sists of  a  mixture  of  hydrocarbons,  the  boiling  points  of 
which  lie  within  certain  limits. 

2.  Coal  tar.  This  product  is  obtained  in  the  manufac- 
ture of  coal  gas,  as  already  explained.  It  is  a  complex 
mixture  and  is  refined  by  the  same  general  method  used 
in  refining  crude  petroleum.  The  principal  hydrocarbons 
obtained  from  the  coal  tar  are  benzene,  toluene,  naphthalene, 
and  anthracene.  In  addition  to  the  hydrocarbons,  coal  tar 
contains  many  other  compounds,  such  as  carbolic  acid  and 
aniline. 

Properties  of  the  hydrocarbons.  The  lower  members  of 
the  first  two  series  of  hydrocarbons  mentioned  are  all 
gases ;  the  succeeding  members  are  liquids.  In  some 
series,  as  the  methane  series,  the  higher  members  are 
solids.  The  preparation  and  properties  of  methane  and 
acetylene  have  been  discussed  in  a  previous  chapter.  Ethy- 
lene  is  present  in  small  quantities  in  coal  gas  and  may  be 


SOME  SIMPLE  ORGANIC  COMPOUNDS         401 

obtained  in  the  laboratory  by  treating  alcohol  (C2H6O) 
with  sulphuric  acid  : 

C2H60  =  C2H4  +  H20. 

Benzene,  the  first  member  of  the  benzene  series,  is  a  liquid 
boiling  at  80°. 

The  hydrocarbons  serve  as  the  materials  from  which  a 
large  number  of  compounds  can  be  prepared  ;  indeed,  it 
has,  been  proposed  to  call  organic  chemistry  the  chemistry 
of  the  hydrocarbon  derivatives. 

Substitution  products  of  the  hydrocarbons.  As  a  rule,  at 
least  a  part  of  the  hydrogen  in  any  hydrocarbon  can  be 
displaced  by  an  equivalent  amount  of  certain  elements  or 
groups  of  elements.  Thus  the  compounds  CH3C1,  CH2C12, 
CHC13,  CC14  can  be  obtained  from  methane  by  treatment 
with  chlorine.  Such  compounds  are  called  substitution 
products. 

Chloroform  (CHC13).  This  can  be  made  by  treating 
methane  with  chlorine,  as  just  indicated,  although  a  much 
easier  method  consists  in  treating  alcohol  or  acetone  (which 
see)  with  bleaching  powder.  Chloroform  is  a  heavy  liquid 
having  a  pleasant  odor  and  a  sweetish  taste.  It  is  largely 
used  as  a  solvent  and  as  an  anaesthetic  in  surgery. 

lodoform  (CHI3).  This  is  a  yellow  crystalline  solid  ob- 
tained by  treating  alcohol  with  iodine  and  an  alkali.  It  has 
a  characteristic  odor  and  is  used  as  an  antiseptic. 

ALCOHOLS 

When  such  a  compound  as  CH3C1  is  treated  with  silver 
hydroxide  the  reaction  expressed  by  the  following  equa- 
tion takes  place  : 

CH3C1  +  AgOH  =  CH3OH  +  AgCl. 


402     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Similarly  C2H5C1  will  give  C2H5OH  and  AgCl.  The  com- 
pounds CH3OH  and  C2H5OH  so  obtained  belong  to  the 
class  of  substances  known  as  alcohols.  From  their  formu- 
las it  will  be  seen  that  they  may  be  regarded  as  derived 
from  hydrocarbons  by  substituting  the  hydroxyl  group 
(OH)  for  hydrogen.  Thus  the  alcohol  CH3OH  may  be 
regarded  as  derived  from  methane  (CH4)  by  substituting 
the  group  OH  for  one  atom  of  hydrogen.  A  great  many 
alcohols  are  known,  and,  like  the  hydrocarbons,  they  may 
be  grouped  into  series.  The  relation  between  the  first 
three  members  of  the  methane  series  and  the  correspond- 
ing alcohols  is  shown  in  the  following  table  : 

CH4  (methane) .     .     ,     .     .    ,,     CH3OH  (methyl  alcohol). 
C2H6  (ethane)    .     .     ...    ..     C2H5OH  (ethyl  alcohol). 

C3H8  (propane)      .     .     '.     .     .     C3H7OH  (propyl  alcohol). 

Methyl  alcohol  (wood  alcohol}  (CH3OH).  When  wood 
is  placed  in  an  air-tight  retort  and  heated,  a  number  of 
compounds  are  evolved,  the  most  important  of  which  are 
the  three  liquids,  methyl  alcohol,  acetic  acid,  and  acetone. 
Methyl  alcohol  is  obtained  entirely  from  this  source,  and 
on  this  account  is  commonly  called  wood  alcohol.  It  is  a 
colorless  liquid  which  has  a  density  of  0.79  and  boils  at  67°. 
It  burns  with  an  almost  colorless  flame  and  is  sometimes 
used  for  heating  purposes,  in  place  of  the  more  expensive 
ethyl  alcohol.  It  is  a  good  solvent  for  organic  substances 
and  is  used  especially  as  a  solvent  in  the  manufacture  of 
varnishes.  It  is  very  poisonous. 

Ethyl  alcohol  (common  alcohol]  (C2H5OH).  i.  Prepara- 
tion. This  compound  may  be  prepared  from  glucose 
(C6H12O6),  a  sugar  easily  obtained  from  starch.  If  some 
baker's  yeast  is  added  to  a  solution  of  glucose  and  the 
temperature  is  maintained  at  about  30°,  bubbles  of  gas  are 


SOME  SIMPLE  ORGANIC  COMPOUNDS 


403 


soon  evolved,  showing  that  a  change  is  taking  place.  The 
yeast  contains  a  large  number  of  minute  organized  bodies, 
which  are  really  forms  of  plant  life.  The  plant  grows  in 
the  glucose  solution,  and  in  so  doing  secretes  a  substance 
known  as  zymase,  which  breaks  down  the  glucose  in  accord- 
ance with  the  following  equation  : 

C6Hi2°6  =  2  C2H5OH  +  2  CO2. 

Laboratory  preparation  of  alcohol.  The  formation  of  alcohol  and 
carbon  dioxide  from  glucose  may  be  shown  as  follows  :  About  loo  g. 
of  glucose  are  dissolved  in  a  liter  of  water  in  flask  A  (Fig.  90). 
This  flask  is  connected  with  the  bottle  B,  which  is  partially  filled 
with  limewater.  The 
tube  C  contains  solid 
sodium  hydroxide.  A 
little  baker's  yeast  is  now 
added  to  the  solution  in 
flask  A,  and  the  appa- 
ratus is  connected,  as 
shown  in  the  figure.  If 
the  temperature  is  main- 
tained at  about  30°,  the 
reaction  soon  begins. 
The  bubbles  of  gas 
escape  through  the  lime- 


FIG.  90 


water  in  B.  A  precipitate  of  calcium  carbonate  soon  forms  in  the 
limewater,  showing  the  presence  of  carbon  dioxide.  The  sodium 
hydroxide  in  tube  C  prevents  the  carbon  dioxide  in  the  air  from 
acting  on  the  limewater.  The  alcohol  remains,  in  the  flask  A  and 
may  be  separated  by  fractional  distillation. 

2.  Properties.  Ethyl  alcohol  is  a  colorless  liquid  with 
a  pleasant  odor.  It  has  a  density  of  0.78  and  boils  at  78°. 
It  resembles  methyl  alcohol  in  its  general  .properties.  It 
is  sometimes  used  as  a  source  of  heat,  since  its  flame  is 
very  hot  and  does  not  deposit  carbon,  as  the  flame  from 
oil  does.  When  taken  into  the  system  in  small  quantities 


404     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

it  causes  intoxication ;  in  larger  quantities  it  acts  as  a 
poison.  The  intoxicating  properties  of  such  liquors  as 
beer,  wine,  and  whisky  are  due  to  the  alcohol  present. 
Beer  contains  from  2  to  5%  of  alcohol,  wine  from  5  to  20%, 
and  whisky  about  50%.  The  ordinary. alcohol  of  the  drug- 
gist contains  94%  of  alcohol  and  6%  of  water.  When  this 
is  boiled  with  lime  and  then  distilled  nearly  all  the  water  is 
removed,  the  distillate  being  called  absolute  alcohol. 

Commercial  preparation  of  alcohol.  Alcohol  is  prepared  commer- 
cially from  starch  obtained  from  corn  or  potatoes.  The  starch  is 
first  converted  into  a  sugar  known  as  maltose,  by  the  action  of  malt, 
a  substance  prepared  by  moistening  barley  with  water,  allowing  it  to 
germinate,  and  then  drying  it.  There  is  present  in  the  malt  a  sub- 
stance known  as  diastase,  which  has  the  property  of  changing  starch 
into  maltose.  This  sugar,  like  glucose,  breaks  down  into  alcohol  and 
carbon  dioxide  in  the  presence  of  yeast.  The  resulting  alcohol  is 
separated  by  fractional  distillation. 

Denatured  alcohol.  The  94%  alcohol  is  prepared  at  present  at  a 
cost  of  about  35  cents  per  gallon,  which  is  about  half  the  cost  of  the 
preparation  of  methyl  alcohol.  The  government,  however,  imposes  a 
tax  on  all  ethyl  alcohol  which  amounts  to  $2.08  per  gallon  on  the 
94%  product.  This  increases  its  cost  to  such  an  extent  that  it  is  not 
economical  to  use  it  for  many  purposes  for  which  it  is  adapted,  such 
as  a  solvent  in  the  preparation  of  paints  and  varnishes  and  as  a 
material  for  the  preparation  of  many  important  organic  compounds. 
By  an  act  of  Congress  in  1906,  the  tax  was  removed  from  denatured 
alcohol,  that  is  alcohol  mixed  with  some  substance  which  renders  it 
unfit  for  the  purposes  of  a  beverage  but  will  not  impair  its  use  for 
manufacturing  purposes.  Some  of  the  European  countries  have  similar 
laws.  The  substances  ordinarily  used  to  denature  alcohol  are  wood  alco- 
hol and  pyridine,  the  latter  compound  having  a  very  offensive  odor. 

Fermentation.  The  reaction  which  takes  place  in  the  preparation 
of  ethyl  alcohol  belongs  to  the  class  of  changes  known  under  the 
general  name  of  fermentation.  Thus  we  say  that  the  yeast  causes 
the  glucose  to  ferment,  and  the  process  is  known  as  alcoholic  fer- 
mentation. There  are  many  kinds  of  fermentations,  and  each  is 
thought  to  be  due  to  the  presence  of  a  definite  substance  known 


SOME  SIMPLE  ORGANIC  COMPOUNDS         405 

as  an  enzyme,  which  acts  by  catalysis.  In  many  cases,  as  in  alco- 
holic fermentation,  the  change  is  brought  about  by  the  action  of 
minute  forms  of  life.  These  probably  secrete  the  enzymes  which 
cause  the  fermentation  to  take  place.  Thus  the  yeast  plant  is  sup- 
posed to  bring  about  alcoholic  fermentation  by  secreting  the  enzyme 
known  as  zymase. 

Glycerin  (C3H5(OH)3).  This  compound  may  be  regarded 
as  derived  from  propane  (C3H8)  by  displacing  three  atoms 
of  hydrogen  by  three  hydroxyl  groups,  and  must  therefore 
be  regarded  as  an  alcohol.  It  is  formed  in  the  manufacture 
of  soaps,  as  will  be  explained  later.  It  is  an  oily,  colorless 
liquid  having  a  sweetish  taste.  It  is  used  in  medicine  and 
in  the  manufacture  of  the  explosives  nitroglycerin  and 
dynamite. 

ALDEHYDES 

When  alcohols  are  treated  with  certain  oxidizing  agents 
two  hydrogen  atoms  are  removed  from  each  molecule  of 
the  alcohol.  The  resulting  compounds  are  known  as  alde- 
hydes. The  relation  of  the  aldehydes  derived  from  methyl 
and  ethyl  alcohol  to  the  alcohols  themselves  may  be  shown 

as  follows  : 

fCH.OH  fCHoO 

'     \C2H5OH  Corresponding  aldehydes    j^^ 

The  first  of  these  (CH2O)  is  a  gas  known  as  formaldehyde. 
Its  aqueous  solution  is  largely  used  as  an  antiseptic  and 
disinfectant  under  the  name  of  formalin.  Acetaldehyde 
(C2H4O)  is  a  liquid  boiling  at  21°. 

ACIDS 

Like  the  other  classes  of  organic  compounds,  the  organic 
acids  may  be  arranged  in  homologous  series.  One  of  the 
most  important  of  these  series  is  the  fatty-acid  series,  the 


406     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

name  having  been  given  to  it  because  the  derivatives  of 
certain  of  its  members  are  constituents  of  the  fats.  Some 
of  the  most  important  members  of  the  series  are  given  in 
the  following  table.  They  are  all  monobasic,  and  this  fact 
is  expressed  in  the  formulas  by  separating  the  replaceable 
hydrogen  atom  from  the  rest  of  the  molecule  : 

H  •  CHO2 formic  acid,  a  liquid  boiling  at  100°. 

H  •  C2H3O2 acetic  acid,  a  liquid  boiling  at  1 18°. 

H  •  C3H5O2 propionic  acid,  a  liquid  boiling  at  140°. 

H-  C4H7O2 butyric  acid,  a  liquid  boiling  at  163°. 

H  •  C16H31O2 palmitic  acid,  a  solid  melting  at  62°. 

H  •  C18H35O2 stearic  acid,  a  solid  melting  at  69°. 

Formic  acid  (H  •  CHO2).  The  name  "  formic  "  is  derived 
from  the  Latin  formica,  signifying  ant.  This  name  was 
given  to  the  acid  because  it  was  formerly  obtained  from  a 
certain  kind  of  ants.  It  is  a  colorless  liquid  and  occurs  in 
many  plants  such  as  the  stinging  nettles.  The  inflamma- 
tion caused  by  the  sting  of  the  bee  is  due  to  formic  acid. 

Acetic  acid  (H-C2H3O2).  Acetic  acid  is  the  acid'pres- 
ent  in  vinegar,  the  sour  taste  being  due  to  it.  It  can  be 
prepared  by  either  of  the  following  methods. 

i.  Acetic  fermentation.  This  consists  in  the  change  of 
'alcohol  into  acetic  acid  through  the  agency  of  a  minute 
organism  commonly  called  mother  of  vinegar.  The  change 
is  represented  by  the  following  equation  : 

C2H5OH  +  2O  =  HC2H3O2  +  H2O. 

The  various  kinds  of  vinegars  are  all  made  by  this  process. 
In  the  manufacture  of  cider  vinegar  the  sugar  present  in 
the  cider  first  undergoes  alcoholic  fermentation ;  the  re- 
sulting alcohol  then  undergoes  acetic  fermentation.  The 
amount  of  acetic  acid  present  in  vinegars  varies  from  3 
to  6%. 


SOME  SIMPLE  ORGANIC  COMPOUNDS         407 

2.  From  the  distillation  of  wood.  The  liquid  obtained 
by  heating  wood  in  the  absence  of  air  contains  a  large 
amount  of  acetic  acid,  and  this  can  be  separated  readily  in 
a  pure  state.  This  is  the  most  economical  method  for  the 
preparation  of  the  concentrated  acid. 

Acetic  acid  is  a  colorless  liquid  and  has  a  strong  pun- 
gent odor.  Many  of  its  salts  are  well-known  compounds. 
Lead  acetate  (Pb(C2H3O2)2)  is  the  ordinary  sugar  of  lead. 
Sodium  acetate  (NaC2H3O2)  is  a  white  solid  largely  used  in 
making  chemical  analyses.  Copper  acetate  (Cu(C2H3O2)2) 
is  a  blue  solid.  When  copper  is  acted  upon  by  acetic  acid 
in  the  presence  of  air  a  green  basic  acetate  of  copper  is 
formed.  This  is  commonly  known  as  verdigris.  All  acetates 
are  soluble  in  water. 

Butyric  acid  (H  •  C4H7O2).  Derivatives  of  butyric  acid 
are  present  in  butter  and  impart  to  it  its  characteristic 
flavor. 

Palmitic  and  stearic  acids.  Ordinary  fats  consist  prin- 
cipally of  derivatives  of  palmitic  and  stearic  acids.  When 
the  fats  are  heated  with  sodium  hydroxide  the  sodium 
salts  of  these  acids  are  formed.  If  hydrpchloric  acid  is 
added  to  a  solution  of  the  sodium  salts,  the  free  palmitic 
and  stearic  acids  are  precipitated.  They  are  white  solids, 
insoluble  in  water.  Stearic  acid  is  often  used  in  making 
candles. 

Acids  belonging  to  other  series.  In  addition  to  members 
of  the  fatty-acid  series,  mention  may  be  made  of  the  fol- 
lowing well-known  acids. 

Oxalic  acid  (H2C2O4).  This  is  a  white  solid  which  occurs 
in  nature  in  many  plants,  such  as  the  sorrels.  Its  ammo- 
nium salt  ((NH4)2C2O4)  is  used  as  a  reagent  for  the  detec- 
tion of  calcium.  When  added  to  a  solution  of  a  calcium 


408     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

compound  the  white,  insoluble  calcium  oxalate  (CaC2O4) 
precipitates. 

Tartaric  acid  (H2  •  C4H4O6).  This  compound  occurs  either 
in  a  free  state  or  in  the  form  of  its  salts  in  many  fruits. 
The  potassium  acid  salt  (KHC4H4O6)  occurs  in  the  juice  of 
grapes.  When  the  juice  ferments  in  the  manufacture  of 
wine,  this  salt,  being  insoluble  in  alcohol,  separates  out  on 
the  sides  of  the  cask  and  in  this  form  is  known  as  argol. 
This  is  more  or  less  colored  by  the  coloring  matter  of  the 
grape.  When  purified  it  forms  a  white  solid  and  is  sold 
under  the  name  of  cream  of  tartar.  The  following  are 
also  well-known  salts  of  tartaric  acid  :  potassium  sodium 
tartrate  (Rochelle  salt)  (KNaC4H4O6),  potassium  antimonyl 
tartrate  (tartar  emetic)  (KSbOC4H4O6). 

Cream  of  tartar  baking  powders.  The  so-called  cream  of  tartar 
baking  powders  consist  of  a  mixture  of  cream  of  tartar,  bicarbonate 
of  soda,  and  some  starch  or  flour.  When  water  is  added  to  this  mix- 
ture the  cream  of  tartar  slowly  acts  upon  the  soda  present  liberating 
carbon  dioxide  in  accordance  with  the  following  equation : 

KHC4H406  +  NaHC03  =  KNaC4H4O6  +  H2O  +  CO2. 

The  carbon  dioxide  evolved  escapes  through  the  dough,  thus  making 
it  light  and  porous. 

Citric  acid  (H3  •  C6H5O7).  This  acid  occurs  in  many 
fruits,  especially  in  lemons.  It  is  a  white  solid,  soluble  in 
water,  and  is  often  used  as  a  substitute  for  lemons  in 
making  lemonade. 

Lactic  acid  (H  •  C3H5O3).  This  is  a  liquid  which  is  formed 
in  the  souring  of  milk. 

Oleic  acid  (H  •  C18H33O2).  The  derivatives  of  this  acid 
constitute  the  principal  part  of  many  oils  and  liquid  fats. 
The  acid  itself  is  an  oily  liquid. 

I 


SOME  SIMPLE  ORGANIC  COMPOUNDS         409 

ETHEREAL   SALTS 

When  acids  are  brought  in  contact  with  alcohols  under 
certain  conditions  a  reaction  takes  place  similar  to  that 
which  takes  place  between  acids  and  bases.  The  following 
equations  will  serve  as  illustrations  : 

KOH  +  HNO3  -  KNO3  +  H2O, 
CH3OH  +  HNO3  =  CH3NO3  +  H2O. 

The  resulting  compounds  of  which  methyl  nitrate  (CH3NO3) 
may  be  taken  as  the  type  belong  to  the  class  known  as 
ethereal  salts,  the  name  having  been  given  them  because 
some  of  them  possess  pleasant  ethereal  odors.  It  will  be 
seen  that  the  ethereal  salts  differ  from  ordinary  salts  in 
that  they  contain  a  hydrocarbon  radical,  such  as  CH3, 
C2H5,  C3H5,  in  place  of  a  metal. 

The  nitrates  of  glycerin  (nitroglycerin}.  Nitric  acid 
reacts  with  glycerin  in  the  same  way  that  it  reacts 
with  a  base  containing  three  hydroxyl  groups  such  as 
Fe(OH)3 : 

Fe(OH)3  +  3  HN03  =  Fe(NO3)3  +  3  H2O, 
C3H6(OH)3  +  3  HN03  =  C8H6(N03)8  +  3  H2O. 

The  resulting  nitrate  (C3H5(NO3)3)  is  the  main  constitu- 
ent of  nitroglycerin^  a  slightly  yellowish  oil  characterized 
by  its  explosive  properties.  Dynamite  consists  of  porous 
earth  which  has  absorbed  nitroglycerin,  and  its  strength 
depends  on  the  amount  present.  It  is  used  much  more 
largely  than  nitroglycerin  itself,  since  it  does  not  explode 
so  readily  by  concussion  and  hence  can  be  transported 
with  safety.' 

The  fats.  These  are  largely  mixtures  of  the  ethereal 
salts  known  respectively  as  olein,  palmitin,  and  stearin. 


410     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

These  salts  may  be  regarded  as  derived  from  oleic,  palmitic, 
and  stearic  acids  respectively,  by  replacing  the  hydrogen  of 
the  acid  with  the  glycerin  radical  C3H5.  Since  this  radical 
is  trivalent  and  oleic,  palmitic,  and  stearic  acids  contain 
only  one  replaceable  hydrogen  atom  to  the  molecule,  it 
is  evident  that  three  molecules  of  each  acid  must  enter 
into  each  molecule  of  the  ethereal  salt.  The  formulas 
for  the  acids  and  the  ethereal  salts  derived  from  each  are 

as  follows  : 

(  HC18H33O2  (oleic  acid) 
I  C3H5(C18H3302)3  (olein) 
f  HC16H31O2  (palmitic  acid) 
I  C3H5(C16H3102)3  (palmitin) 
<"  HC18H35O2  (stearic  acid) 
I  C3H5(C18H3502)3  (stearin) 

Olein  is  a  liquid  and  is  the  main  constituent  of  liquid 
fats.  Palmitin  and  stearin  are  solids. 

Butter  fat  and  oleomargarine.  Butter  fat  consists  prin- 
cipally of  olein,  palmitin,  and  stearin.  The  flavor  of  the 
fat  is  due  to  the  presence  of  a  small  amount  of  butyrin, 
which  is  an  ethereal  salt  of  butyric  acid.  Oleomargarine 
differs  from  butter  mainly  in  the  fact  that  a  smaller 
amount  of  butyrin  is  present.  It  is  made  from  the  fats 
obtained  from  cattle  and  hogs.  This  fat  is  churned  up 
with  milk,  or  a  small  amount  of  butter  is  added,  in  order  to 
furnish  sufficient  butyrin  to  impart  the  butter  flavor. 

Saponification.  When  an  ethereal  salt  is  heated  with  an 
alkali  a  reaction  expressed  by  the  following  equation  takes 

place:     C2H5NO3  +  KOH  =  C2H5OH  +  KNO3. 

This  process  is  known  as  saponification,  since  it  is  the 
one  which  takes  place  in  the  manufacture  of  soaps.  The 
ordinary  soaps  are  made  by  heating  fats  with  a  solution  of 


SOME  SIMPLE  ORGANIC  COMPOUNDS         411 

sodium  hydroxide.  The  reactions  involved  may  be  illus- 
trated by  the  following  equation  representing  the  reaction 
between  palmitin  and  sodium  hydroxide  : 

C8H6(C16H8102)8  +  3  NaOH  =  3  NaC16H31O2  + 
C3H6(OH)3. 

In  accordance  with  this  equation  the  ethereal  salts  in  the 
fats  are  converted  into  glycerin  and  the  sodium  salts  of  the 
corresponding  acids.  The  sodium  salts  are  separated  and 
constitute  the  soaps.  These  salts  are  soluble  in  water. 
When  added  to  water  containing  calcium  salts  the  insol- 
uble calcium  palmitate  and  stearate  are  precipitated.  Mag- 
nesium salts  act  in  a  similar  way.  It  is  because  of  these 
facts  that  soap  is  used  up  by  hard  waters. 

ETHERS 

When  ethyl  alcohol  is  heated  to  140°  with  sulphuric 
acid  the  reaction  expressed  by  the  following  equation  takes 

place  :  2  C2H5OH  =  (C2H6)20  +  H20. 

The  resulting  compound,  (C2H5)2O,  is  ordinary  ether  and 
is  the  most  important  member  of  the  class  of  compounds 
called  ethers.  Ordinarily  ether  is  a  light,  very  inflammable 
liquid  boiling  at  35°.  It  is  used  as  a  solvent  for  organic 
substances  and  as  an  anaesthetic  in  surgical  operations. 

KETONES 

The  most  common  member  of  this  group  is  acetone 
(C3H6O),  a  colorless  liquid  obtained  when  wood  is  heated 
in  the  absence  of  air.  It  is  used  in  the  preparation  of 
other  organic  compounds,  especially  chloroform. 


412     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

ORGANIC  BASES 

This  group  includes  a  number  of  compounds,  all  of  which 
contain  nitrogen  as  well  as  carbon.  They  are  characterized 
by  combining  directly  with  acids  to  form  salts,  and  in  this 
respect  they  resemble  ammonia.  They  may,  indeed,  be 
regarded  as  derived  from  ammonia  by  displacing  a  part  or 
all  of  the  hydrogen  present  in  ammonia  by  hydrocarbon 
radicals.  Among  the  simplest  of  these  compounds  may 
be  mentioned  methylamine  (CH3NH2)  and  ethylamine 
(C2H5NH2).  These  two  compounds  are  gases  and  are 
formed  in  the  distillation  of  wood  and  bones.  Pyridine 
(C5H5N)  and  quinoline  (C9H7N)  are  liquids  present  in 
small  amounts  in  coal  tar,  and  also  in  the  liquid  obtained 
by  the  distillation  of  bones.  Most  of  the  compounds  now 
classified  under  the  general  name  of  alkaloids  (which  see) 
also  belong  to  this  group. 

CARBOHYDRATES 

s 

The  term  ''carbohydrate  "  is  applied  to  a  class  of  com- 
pounds which  includes  the  sugars,  starch,  and  allied  bodies. 
These  compounds  contain  carbon,  hydrogen,  and  oxygen, 
the  last  two  elements  generally  being  present  in  the  pro- 
portion in  which  they  combine  to  form  water.  The  most 
important  members  of  this  class  are  the  following: 

Cane  sugar C12H22O11. 

Milk  sugar ^12^22^11- 

Glucose C6H12O6. 

Levulose C6H12O6. 

Cellulose C6H10O5. 

Starch  .     . C6H1005. 

Cane  sugar  (C12H22On).  This  is  the  well-known  sub- 
stance commonly  called  sugar.  It  occurs  in  many  plants, 


SOME  SIMPLE  ORGANIC  COMPOUNDS         413 

especially  in  the  sugar  cane  and  sugar  beet.  It  was  for- 
merly obtained  almost  entirely  from  the  sugar  cane,  but  at 
present  the  greatest  amount  of  it  comes  from  the  sugar 
beet.  The  juice  from  the  cane  or  beet  contains  the  sugar 
in  solution  along  with  many  impurities.  These  impurities 
are  removed,  and  the  resulting  solution  is  then  evaporated 
until  the  sugar  crystallizes  out.  The t  evaporation  is  con- 
ducted in  closed  vessels  from  which  the  air  is  partially 
exhausted.  In  this  way  the  boiling  point  of  the  solution 
is  lowered  and  the  charring  of  the  sugar  is  prevented. 
It  is  impossible  to  remove  all  the  sugar  from  the  solution. 
In  preparing  sugar  from  sugar  cane  the  liquors  left  after 
separating  as  much  of  it  as  possible  from  the  juice  of  the 
cane  constitute  ordinary  molasses.  Maple  sugar  is  made  by 
the  evaporation  of  the  sap  obtained  from  a  species  of  the 
maple  tree.  Its  sweetness  is  due  to  the  presence  of  cane 
sugar,  other  products  present  in  the  maple  sap  imparting 
the  distinctive  flavor. 

When  a  solution  of  cane  sugar  is  heated  with  hydro- 
chloric or  other  dilute  mineral  acid,  two  compounds,  glu- 
cose and  levulose,"  are  formed  in  accordance  with  the 
following  equation : 

^12^22^11  -f  H2O  =  C6H12O6  +  C6H12O6. 

This  same  change  is  brought  about  by  the  action  of  an 
enzyme  present  in  the  yeast  plant.  When  yeast  is  added  to 
a  solution  of  cane  sugar  fermentation  is  set  up.  The  cane 
sugar,  however,  does  not  ferment  directly :  the  enzyme  in 
the  yeast  first  transforms  the  sugar  into  glucose  and  levu- 
lose, and  these  sugars  then  undergo  alcoholic  fermentation. 
When  heated  to  1 60°  cane  sugar  melts ;  if  the  temper- 
ature is  increased  to  about  215°,  a  partial  decomposition 


4 14     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

takes  place  and  a  brown  substance  known  as  caramel  forms. 
This  is  used  largely  as  a  coloring  matter. 

Milk  sugar  (C12H22On).  This  sugar  is  present  in  the 
milk  of  all  mammals.  The  average  composition  of  cow's 
milk  is  as  follows  : 

Water 87.17% 

Casein  (nitrogenous  matter)    .     .     .     .       3.56 

Butter  fat 3.64 

Sugar  of  milk 4.88 

Mineral  matter 0.75 

When  rennin,  an  enzyme  obtained  from  the  stomach  of 
calves,  is  added  to  milk,  the  casein  separates  and  is  used 
in  the  manufacture  of  cheese.  The  remaining  liquid  con- 
tains the  milk  sugar  which  separates  on  evaporation ;  it 
resembles  cane  sugar  in  appearance  but  is  not  so  sweet 
or  soluble.  The  souring  of  milk  is  due  to  the  fact  that 
the  milk  sugar  present  undergoes  lactic  fermentation  in 
accordance  with  the  equation 

C12H22On  +  H20  =  4  C3H603. 

The  lactic  acid  formed  causes  the  separation  of  the  casein, 
thus  giving  the  well-known  appearance  of  sour  milk. 

Isomeric  compounds.  It  will  be  observed  that  cane  sugar  and  milk 
sugar  have  the  same  formulas.  Their  difference  in  properties  is  due 
to  the  different  arrangement  of  the  atoms  in  the  molecule.  Such 
compounds  are  said  to  be  isomeric.  Glucose  and  levulose  are  also 
isomeric. 

Glucose  (grape  sugar,  dextrose]  (C6H12O6).  This  sugar 
is  present  in  many  fruits  and  is  commonly  called  grape 
sugar  because  of  its  presence  in  grape  juice.  It  can  be 
obtained  by  heating  cane  sugar  with  dilute  acids,  as 


SOME  SIMPLE  ORGANIC  COMPOUNDS         415 

explained  above  ;  also  by  heating  starch  with  dilute  acids, 
the  change  being  as  follows  : 

C6H10O5  +  H2O  =  C6H12O6. 

Pure  glucose  is  a  white  crystalline  solid,  readily  soluble  in 
water,  and  is  not  so  sweet  as  cane  sugar.  In  the  presence 
of  yeast  it  undergoes  alcoholic  fermentation.  It  is  pre- 
pared from  starch  in  large  quantities,  and  being  less  expen- 
sive than  cane  sugar,  is  used  as  a  substitute  for  it  in  the 
manufacture  of  jellies,  jams,  molasses,  candy,  and  other 
sweets.  As  commonly  sold  on  the  market  it  is  in  the 
form  of  a  thick  colorless  sirup. 

Levulose  (fruit  sugar}  (C6H12O6).  This  sugar  is  a  white 
solid  which  occurs  along  with  glucose  in  fruits  and  honey. 
It  undergoes  alcoholic  fermentation  in  the  presence  of 
yeast. 

Cellulose  (C6H10O5).  This  forms  the  basis  of  all  woody 
fibers.  Cotton  and  linen  are  nearly  pure  cellulose.  It  is 
insoluble  in  water,  alcohol,  and  dilute  acids.  Sulphuric  acid 
slowly  converts  it  into  glucose.  Nitric  acid  forms  nitrates 
similar  to  nitroglycerin  in  composition  and  explosive 
properties.  These  nitrates  are  variously  known  as  nitro- 
cellulose, pyroxylin,  and  gun  cotton.  When  exploded  they 
yield  only  colorless  gases  ;  hence  they  are  used  especially 
in  the  manufacture  of  smokeless  gunpowder.  Collodion 
is  a  solution  of  nitrocellulose  in  a  mixture  of  alcohol  and 
ether.  Celluloid  is  a  mixture  of  nitrocellulose  and  cam- 
phor. Paper  consists  mainly  of  cellulose,  the  finer  grades 
being  made  from  linen  and  cotton  rags,  and  the  cheaper 
grades  from  straw  and  wood. 

Starch  (C6H10O5).  This  is  by  far  the  most  abundant 
carbohydrate  found  in  nature,  being  present  especially  in 


41 6     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

seeds  and  tubers.  In  the  United  States  it  is  obtained 
chiefly  from  corn,  nearly  80%  of  which  is  starch.  In 
Europe  it  is  obtained  principally  from  the  potato.  It  con- 
sists of  minute  granules  and  is  practically  insoluble  in 
cold  water.  These  granules  differ  somewhat  in  appear- 
ance, according  to  the  source  of  the  starch,  so  that  it  is 
often  possible  to  determine  from  what  plant  the  starch  was 
obtained.  When  heated  with  water  the  granules  burst  and 
the  starch  partially  dissolves.  Dilute  acids,  as  well  as  certain 
enzymes,  convert  it  into  glucose  or  similar  sugars.  When 
seeds  germinate  the  starch  present  is  converted  into  solu- 
ble sugars,  which  are  used  as  food  for  the  growing  plant. 

Chemical  changes  in  bread  making.  The  average  com- 
position of  wheat  flour  is  as  follows  : 

Water.     .     .     .     .     .     .".     .     .   :.     .  13.8% 

Protein  (nitrogenous  matter)    ....  7.9 

Fats 1.4 

Starch      .     .     .     .     .:    .     »    ...    .     .     .  76.4 

Mineral  matter 0.5 

In  making  bread  the  flour  is  mixed  with  water  and 
yeast,  and  the  resulting  dough  set  aside  in  a  warm  place 
for  a  few  hours.  The  yeast  first  converts  a  portion  of  the 
starch  into  glucose  or  a  similar  sugar,  which  then  under- 
goes alcoholic  fermentation.  The  carbon  dioxide  formed 
escapes  through  the  dough,  making  it  light  and  porous. 
The  yeast  plant  thrives  best  at  about  30° ;  hence  the 
necessity  for  having  the  dough  in  a  warm  place.  If  the 
temperature  rises  above  50°,  the  vitality  of  the  yeast  is 
destroyed  and  fermentation  ceases.  In  baking  the  bread, 
the  heat  expels  the  alcohol  and  also  expands  the  bubbles 
of  carbon  dioxide  caught  in  the  dough,  thus  increasing  its 
lightness. 


SOME  SIMPLE  ORGANIC  COMPOUNDS         417 

SOME  DERIVATIVES  OF  BENZENE 

Attention  has  been  called  to  the  complex  nature  of 
coal  tar.  Among  the  compounds  present  are  the  hydro- 
carbons, benzene,  toluene,  naphthalene,  and  anthracene. 
These  compounds  are  not  only  useful  in  themselves  but 
serve  for  the  preparation  of  many  other  important  com- 
pounds known  under  the  general  name  of  coal-tar  products. 

Nitrobenzene  (oil of  myrbane)  (C6H5NO2).  When  benzene 
is  treated  with  nitric  acid  a  reaction  takes  place  which  is 
expressed  by  the  following  equation  : 

C6H6  +  HN03  =  C6H6N02  +  H2O. 

The  product  C6H5NO2  is  called  nitrobenzene.  It  is  a 
slightly  yellowish  poisonous  liquid,  with  a  characteristic 
odor.  Its  main  use  is  in  the  manufacture  of  aniline. 

Aniline  (C6H5NH2).  When  nitrobenzene  is  heated  with 
iron  and  hydrochloric  acid  the  hydrogen  evolved  by  the 
action  of  the  iron  upon  the  acid  reduces  the  nitrobenzene 
in  accordance  with  the  following  equation  : 

C6H5N02  +  6  H  =  C6H5NH2  +  2  H2O. 

The  resulting  compound  is  known  as  aniline,  a  liquid 
boiling  at  182°.  When  first  prepared  it  is  colorless,  but 
darkens  on  standing.  Large  quantities  of  it  are  used  in 
the  manufacture  of  the  aniline  or  coal-tar  dyes,  which 
include  many  important  compounds. 

Carbolic  acid  (C6H5OH).  This  compound,  sometimes 
known  as  phenol,  occurs  in  coal  tar,  and  is  also  prepared 
from  benzene.  It  forms  colorless  crystals  which  are 
very  soluble  in  water.  It  is  strongly  corrosive  and  very 
poisonous. 


41 8     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Naphthalene  and  anthracene.  These  are  hydrocarbons 
occurring  along  with  benzene  in  coal  tar.  They  are  white 
solids,  insoluble  in  water.  The  well-known  moth  balls  are 
made  of  naphthalene.  Large  quantities  of  naphthalene  are 
used  in  the  preparation  of  indigo,  a  dye  formerly  obtained 
from  the  indigo  plant,  but  now  largely  prepared  by  labora- 
tory methods.  Similarly  anthracene  is  used  in  the  prepa- 
ration of  the  dye  alizarin,  which  was  formerly  obtained 
from  the  madder  root. 

THE  ALKALOIDS 

This  term  is  applied  to  a  group  of  compounds  found  in 
many  plants  and  trees.  They  all  contain  nitrogen,  and 
most  of  them  are  characterized  by  their  power  to  combine 
with  acids  to  form  salts.  This  property  is  indicated 
by  the  name  alkaloids,  which  signifies  alkali-like.  The 
salts  are  soluble  in  water,  and  on  this  account  are  more 
largely  used  than  the  free  alkaloids,  which  are  insoluble  in 
water.  Many  of  the  alkaloids  are  used  in  medicine,  some 
of  the  more  important  ones  being  given  below. 

Quinine.  This  alkaloid  occurs  along  with  a  number  of 
others  in  the  bark  of  certain  trees  which  grow  in  districts 
in  South  America  and  also  in  Java  and  other  tropical 
islands.  It  is  a  white  solid,  and  its  sulphate  is  used  in 
medicine  in  the  treatment  of  fevers. 

Morphine.  When  incisions  are  made  in  the  unripe  cap- 
sules of  one  of  the  varieties  of  the  poppy  plant,  a  milky 
juice  exudes  which  soon  thickens.  This  is  removed  and 
partially  dried.  The  resulting  substance  is  the  ordinary 
opium  which  contains  a  number  of  alkaloids,  the  principal 
one  being  morphine.  This  alkaloid  is  a  white  solid  and 
is  of  great  service  in  medicine. 


SOME  SIMPLE  ORGANIC  COMPOUNDS         419 

Among  the  other  alkaloids  may  be  mentioned  the  fol- 
lowing :  Nicotine,  a  very  poisonous  liquid,  the  salts  of 
which  occur  in  the  leaves  of  the  tobacco  plant ;  cocaine,  a 
crystalline  solid  present  in  coca  leaves  and  used  in  medicine 
as  a  local  anaesthetic ;  atropine,  a  solid  present  in  the 
berry  of  the  deadly  nightshade,  and  used  in  the  treatment 
of  diseases  of  the  eye ;  strychnine,  a  white,  intensely 
poisonous  solid  present  in  the  seeds  of  the  members  of 
the  Strychnos  family. 


INDEX 


Acetaldehyde  .  . 
Acetic  acid  ,  .  . 
Acetone  .... 
Acetylene  .  .  . 

series  .  .  . 
Acids 

binary    .     .     . 

characteristics 

definition    .     . 

dibasic  .     .     . 

familiar      .     . 

monobasic 

nomenclature 


organic 

preparation 

strength 

ternary  

undissociated  .  .  .  . 

Acker  furnace 

Agate 

Air 

a  mechanical  mixture  .     . 

carbon  dioxide  in    .     .     . 

changes  in  composition    . 

liquid 

nitrogen  in 

oxygen  in 

poisonous  effects  of  ex- 
haled   

properties 

quantitative  analysis  of    . 

regarded  as  an  Clement    . 

standard  for  density    .     . 

water  vapor  in  .  .  .  . 

Alabaster 

Alchemists 

Alchemy 

Alcohol,  common  .  .  .  . 

denatured        

ethyl 

methyl 

wood  ,     . 


PAGE  PAGE 

405  Alcohols 401 

406  Aldehydes 405 

411       Alizarin 418 

203       Alkali 107,  274 

399              family 274 

1 06  Alkaline-earth  family      .     .     .  300 

113       Alkaloids 418 

106  Allotropic  forms 22 

107  Alloys 252 

159      Alum 333 

1 06  ammonium 334 

159  ammonium  chrome      .     .  384 

113  ammonium  iron       .     ,     .  352 

405  baking  powders  ....  335 

141               potassium 333 

1 1 1  potassium  chrome  ...  384 

113  potassium  iron    ....  352 

107  Aluminates 332 

279       Aluminium 327 

260  bronze 33°»359 

83              chloride 333 

89  family 327 

87               hydroxide 332 

87               metallurgy 328 

91               occurrence 327 

87  oxide 331 

85               preparation 328 

properties 329 

88  silicates 335 

90  uses 330 

85       Amalgam 362 

83       Amethyst 260,331 

229       Ammonia 123 

87               composition 127 

308               preparation 123 

9               properties 124 

9              uses 125 

402       Ammonium 126 

404  acid  carbonate    ....  295 

402               carbonate 295 

402               chloride 294 

402              compounds 294 

421 


422      AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


'35 

243 


PAGE 

Ammonium  hydrosulphide      .  296 

hydroxide 126 

molybdate 388 

oxalate 407 

sulphate 295 

sulphide 295 

sulphide,  yellow  ....  296 

Analysis 40 

Anhydride '.  135 

carbonic 206 

chromic 387 

nitric      . 135 

nitrous 

phosphoric      .     .     . 
sulphuric    .... 

Anhydrite 

Aniline 417 

Anion 106 

Anode 99 

Anthracene 418 

Antimony 250 

acids      251 

alloys 253 

chloride 252 

metallic  properties  .     .     .  252 

occurrence 251 

oxides 251 

preparation 251 

properties  .     .     .     .  '  .     .  251 

sulphides 251 

Apatite     .....     175,239,311 

Aqua  ammonia 124 

Aqua  regia 185 

Aqueous  tension 25 

Argon 80 

Arsenic 246 

acids       .......  250 

antidote 250 

Marsh's  test 248 

occurrence 246 

oxides 249 

preparation 246 

properties 247 

sulphides 250 

white 249 

Arsenopyrites 246 

Arsine 247 

Asbestos 321,  336 

Atmosphere 83 

constituents    ......  83* 

function  of  constituents  .  84 


Atomic  hypothesis  .... 
theory  

and  laws  of  matter    . 

and  radium  .  .  . 
weights 

accurate  determina- 
tion   

and  general  proper- 
ties   

and  specific  heats 

calculation  of  ... 

Dalton's  method  .     . 

direct  determination  . 

from  molecular 
weights  .... 

relation  to  equivalent 

standard  for     .     .     . 

steps  in  determining  . 


Atoms 

size 

Atropine 

Aurates 

Avogadro's  hypothesis   .     .     . 

and   chemical    calcula- 
tions   

and  molecular  weights 

Azote . 

Azurite 

Babbitt  metal 

Bacteria 

decomposition    of   organic 
matter  by 

nitrifying 

Baking  powders     .     .     .     .285, 

alum 

soda 

Barium 

chloride 

nitrate 

oxides    .     .     

sulphate 

Barytes 

Bases 

characteristics     .... 

definition 

familiar 

nomenclature 


PAGE 
61 
59 
63 

°65 
231 

167 

233 
231 

223 
233 

230 
224 

66 
224 

62 

65 
419 

396 
226 

235 
227 

78 

357 

253 
85 

122 

85 

408 

335 
312 

3M 

712 


organic  .     . 

strength 

undissociated 


107 
107 
1 08 
107 

"3 

412 

"3 

1 08 


INDEX 


423 


PAGE 

Basic  lining  process  ....  346 

Bauxite ^~~  332 

Beer 404 

Benzene 417 

derivatives 417 

series '.     .  399 

Benzine 400 

Bessemer  process       ....  345 

Bismuth 253 

basic  salts 255 

chloride 253 

nitrate 253 

occurrence 253 

oxides 254 

preparation 253 

salts,  hydrolysis  of .     .     .  254 

subnitrate 256 

uses 253 

Bismuthyl  chloride    ....  256 

Blast  furnace 341 

lamp 38 

Bleaching  powder      ....  306 

Bleaching  by  chlorine     .     .     .  181 

by  sulphurous  acid .     .     .  152 

Boiler  scale  .     .     .    t.     .     .     .  320 

Bone  ash.     .     .     .     .     .     .     .  311 

Bone  black 200 

Borax 265 

bead 266 

Bornite 357 

Boron.     .     . 257,264 

acids 265 

fluoride 264 

hydride 264 

occurrence 264 

oxides 264 

preparation 264 

properties 264 

Brass 323 

Bread  making 416 

Bromides 190 

Bromine 187 

occurrence 187 

oxygen  compounds      .     .  190 

preparation 187 

properties  .     .     .     .     .     .  188 

Bronze 359 

aluminium      ....  330,  359 

Butter  fat 410 

Butyric  acid 407 

By-product 284 


PAGE 

Cadmium 325 

compounds 326 

Caesium 294 

Calamine 321 

Calcite 305 

Calcium 301 

carbide  ......  203,  310 

carbonate 305 

chloride 306 

fluoride .......  308 

hydroxide 303 

occurrence 301 

oxide 302 

phosphate 246,311 

preparation 302 

sulphate 308 

Calomel 363 

Calorie 76 

Caramel 414 

Carbohydrates 413 

Carbolic  acid 417 

Carbon 196 

allotropic  forms  ....  196 

•amorphous 198 


crystalline  forms 
cycle  in  nature    .     . 
dioxide  
and  bases     .     . 
and  plant  life    . 
in  air  . 

*y~ 

•     •      197 
.    .      88 
.    .     204 
.    .     206 
.     .      88' 
87 

occurrence  .     . 
preparation  . 
properties    .     . 
solid   .... 
disulphide  .... 
familv    . 

.     .     204 
.    .     204 
.     .    204 
.    .     204 
.  160,  210 

1  06 

hydrogen  compounds  .     .  201 

monoxide 208 

occurrence 196 

oxides 203 

properties 200 

pure 198 

retort 199 

uses 200 

Carbonates 207 

acid 207 

Carbonic  acid 206 

Carborundum 259 

Carnallite     .     .  % 288 

Casein _.    .     .  414 


424     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


: 

PAGE 

\ 

PAGE 

Cassiterite    

370 

Chromite      

383 

Catalysis  

Chromium    

383 

Catalyzers    

153 

a  base-forming  element    . 

383 

Cathode  

99 

an  acid-forming  element  . 

38S 

Cation      

1  06 

occurrence      

383 

Caustic  potash  

288 

Cinnabar  

363 

soda  

278 

Citric  acid    

408 

Celestite  

312 

Clay     

336 

Celluloid  

Coal     

199 

Cellulose.    

4J5 

gas    

217 

Cement    

304 

products     

400 

Ceramic  industries     .... 

336 

tar     

218 

Cerium     

377 

Cobalt      

3S4 

Chalcedony  

260 

compounds     .     .     .     .     . 

354 

Chalcocite    

3S7 

Cocaine    

419 

Chalcopyrite      

357 

Coke    

199 

Chalk  

30  s 

Collodion      

4*5 

Chamber  acid   

Colemanite  

26s 

Changes,physical  andchemical 

2 

Combining  weights    .... 

225 

Charcoal  

199 

Combustion       

Chemical  affinity  

12 

/broad  sense    

20 

changes      

2 

/  in  air      

19 

compounds     

7 

1  phlogiston  theory    .     .     . 

19 

equilibrium     

128 

products     

18 

properties  

3 

1   spontaneous   

20 

Chemistry,  definition      .     .     . 

4 

1  supporters'      

213 

Chili  saltpeter  191, 

28S 

Compounds,  chemical    .     .     . 

7 

China\Jare    

336 

\isomeric     

414 

Chloria  acid      

187 

^f  metals,  preparation 

26s 

Chlorides      

1  86 

structure  of    

118 

Chlorine  

177 

Conservation  of  energy  .     .     . 

4 

bleaching  action      .     .     . 

181 

of  matter   

S 

chemical  properties      .     . 

180 

Contact  process     

iS4 

family    

174 

Converter,  Bessemer      .     .     . 

345 

historical    

177 

Copper     

357 

occurrence      

178 

acetate  

407 

oxides    

187 

alloys  of     

3S9 

oxygen  acids  
preparation     

187 
178 

family    
hydroxide  

356 
360 

properties  

179 

metallurgy      

357 

Chloroform  

401 

occurrence      

3S7 

Chloroplatinic  acid    .... 
Chlorous  acid   

393 

187 

ores  
oxide     

£ 

Chromates    

38S 

properties  

3S8 

Chrome  alum    

384 

refining  

3S8 

Chromic  acid    

388 

sulphate     

36i 

anhydride  

387 

sulphide     

361 

chloride      

383 

uses  

359 

hydroxide  

383 

Copperas      .     .     ..    .     .     .     . 

35o 

sulphate     

384 

Coral  

30  s 

384 

Corrosive  sublimate  .... 

363 

INDEX 


425 


PAGE 

Corundum 331 

Cream  of  tartar 408 

Crocoisite 383 

Cryolite 175,328 

Crystallization 98 

water  of 54>  75 

Crystallography 161 

Crystals 161 

axes  of 161 

systems 162 

Cupric  compounds     ....  360 

Cuprite 360 

Cuprous  compounds  ....  360 

chloride 360 

oxide 360 

Cyanides 210 

solutions  are  alkaline  .     .  210 

Dalton's  atomic  hypothesis     .  61 

Decay 21 

Decomposition  of  organic  mat- 
ter       122 

Decrepitation 55 

Deliquescence 55 

Density  of  gases 230 

Desiccating  agents     ....  55 

Developers 367 

Dewar  bulb 91 

Dextrose 414 

Diamond 197 

Dichromates      .     .     .     .  r  .     .  385 

Dichromic  acid 385 

Dimorphous  substances      .     .  163 

Dissociation 99 

and  boiling  point     .     .     .  101 

and  freezing  point  .     .     .  101 

equations  of 112 

extent  of 113 

Distillation 50 

Dogtooth  spar 306 

Dolomite 319 

Double  decomposition   ...  71 

Drummond  light 38 

Dyeing 333 

Dynamite 409 

Earth  metals 327 

Efflorescence 54 

Electric  furnace 221 

Electro-chemical  industries      .  269 

Electrode 99 


PAGE 

Electrolysis 99 

of  sodium  chloride  .     .     .  102 

of  sodium  sulphate      .     .  103 

of  water 41,  102 

Electrolytes 99 

Electrolytic  dissociation      .     .  99 

Electroplating 366 

Electrotyping 359 

Elements,  definition  ....  8 

atomic  weights   .     .     ...  232 

earlier  classification     .     .  165 

names 1 1 

natural  groups     ....  165 

number  of 9 

occurrence 10 

periodic  division       .     .     .  166 

physical  state      ....  10 

symbols  of n 

Emery 331 

Energy 4 

and  plant  life       ....  89 

chemical 5 

conservation  of  ....  4 

transformation  of    ...  5 

Enzyme 405 

Epsom  salts 320 

Equations 68 

are  quantitative  ....  72 

knowledge  requisite  for   .  69 

not  algebraic 74 

reading  of 69 

Equilibrium 138 

chemical 138 

in  solution 139 

point  of 138 

Equivalent 224 

determination  of      ...  224 
elements  with   more   than 

one 225 

relation  to  atomic  weight  .  224 

Etching 177 

Ether ;  411 

Ethereal  salts 409 

Ethers 411 

Ethylamine 412 

Ethylene  series 399 

Eudiometer 43 

Evaporation n 

Families  in  periodic  groups     .  170 

triads     . 165 


426     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


PAGE 

Family  resemblances      .     .     .  170 

Fats 409 

Fatty  acid  series 405 

Feldspar 261,335 

Fermentation 404 

acetic 406 

alcoholic 404,  405 

lactic 414 

Ferric  chloride 352 

hydroxide 352 

salts 351 

reduction 353 

sulphate 352 

Ferrochromium      .....  383 

Ferromanganese 343 

Ferrosilicon 259 

Ferrous  carbonate      .     .     .     .  351 

salts 350 

oxidation  of     ....  353 

sulphate 350 

sulphide 350 

Fertilizers 245 

Filtration 6,  51 

beds 52 

Fire  damp 202 

Flames 213 

appearance 214 

blowpipe 216 

Bunsen 214 

conditions  for     .     .     .     .  213 

hydrogen 34 

luminosity 216 

oxidizing 214 

oxyhydrogen 37 

reactions    .     .     .     .     .     .  296 

reducing 214 

structure 214 

Flash  lights 317 

Flint 260 

Fluorides 177 

Fluorine 175 

Fluorspar 175,308 

Fluosilicic  acid 259 

Flux 340 

Fool's  gold 351 

Formaldehyde •    .  405 

Formalin 405 

Formic  acid 406 

Formulas '68 

how  determined ....  234 

structural  .     .     .     .     .     .  119 


PAGE 

Fractional  distillation     ...  51 

Franklinite 321 

Fuels 220 

Furnace,  arc 221 

electric 221 

resistance  , 221 

Fusion  methods 271 

Galena 373 

Gallium 327 

Galvanized  iron     .     .     .     .     .  323 

Gas,  collection  of 15 

coal 217 

fuel 217 

illuminating 217 

measurement  of ....  23 

natural 219 

purification  of     .     .     .     .  218 

water 219 

Gases,  table 220 

Gasoline 400 

German  silver 323,  359 

Germanium 370 

Germs,  effect  of  cold  on      .     .  53 

in  air 84 

in  water 52 

Glass— T* 262 

Xcoloring  of 263 

etching  of 177 

molding  of 263 

nature  of 263 

varieties 263 

Glauber's  salt 281 

Glazing .     .  336 

Glucose 414 

Glycerin 405 

nitrates  of 409 

Gold 393 

alloys 396 

chloride 396 

coin  . 359 

extraction  of 394 

in  copper 358 

mining 394 

occurrence 393 

properties 396 

refining  of 395 

telluride 394 

Goldschmidt  method      .     .  269, 330 

Gram-molecular  weight  .     .     .  236 

Granite 336 


INDEX 


427 


PAGE 

Graphite 198 

Gun  cotton 415 

metal 359 

powder 292 

Gypsite 308 

Gypsum 308 

Halogens 174 

Hard  water 309 

Heat  of  reaction 75 

Helium 80,314 

Hematite 339,  349 

Homologous  series    ....  398 

Hydriodic  acid 193 

Hydrobromic  acid      ....  189 

Hydrocarbons 201,  398 

properties 400 

series 398 

substitution  products  .     .  401 

Hydrochloric  acid      ....  182 

composition 183 

oxidation  of 185 

preparation 182 

properties 184 

salts 186 

Hydrocyanic  acid 210 

Hydrofluoric  acid 176 

etching  by 177 

salts  of .     .     .     .     .     .     .  177 

Hydrogen 28 

dioxide 56 

explosive  with  oxygen      .  35 

occurrence 28 

preparation  from  acids     .  30 

preparation  from  water    .  28 

properties 32 

standard       for       atomic 

weights 66 

standard    for    molecular 

weights 227 

sulphide     ......  146 

uses .  38 

Hydrolysis 254 

conditions  affecting     .     .  255 

partial 255 

Hydrosulphuric  acid ....  1 46 

Hydroxyl  radical 112 

Hypochlorous  acid    .     .     .     .  187 

Hypothesis 61 

Avogadro's 226 

Dalton's 61 


PAGE 

Ice  manufacture 125 

Iceland  spar 305 

Indigo 418 

Indium 327 

Insoluble  compounds     .     .     .  272 

lodic  acid     . 194 

Iodides 193 

lodme 190 

oxygen  compounds      .     .  193 

preparation 191 

properties 192 

tincture 192 

lodoform 192,401 

Ions 100 

and  electrolytes  ....  104 

Iridium 393 

Iron 339 

alum 352 

cast -..  343 

compounds 349 

cyanides 352 

disulphide 351 

family 338 

metallurgy 339 

occurrence 339 

ores 339 

oxides 349 

pure 348 

varieties 342,  347 

wrought 343 


Jasper 


260 


__  „-     261,335 

Kerosene 400 

Ketones 411 

Kieserite 288 

Kindling  temperature     ...  17 

Krypton 80 

Lactic  acid 408 

Lampblack 200 

Laughing  gas 132 

Law,  definition 61 

of  Boyle 24 

of  Charles 23 

of  combining  volumes      .  194 

of  conservation  of  energy  4 

of  conservation  of  matter  5,  59 

of  definite  composition    .  59 


428      AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


PAGE 

Law  of  Dulorig  and  Petit  .     .  233 

of  Gay-Lussac     .     .  j  .     .  194 

of  'multiple  proportion      .  60 

of  Raoult 233 

periodic 169 

Lead 373 

acetate 375,  407 

alloys 375 

basic  carbonate  ....  376 

carbonate 376 

chloride 377 

chromate 377 

insoluble  compounds  .     .  376 

metallurgy 373 

nitrate 375 

occurrence 373 

oxides 375 

peroxide 375 

properties 374 

red 375 

soluble  salts 375 

sugar  of 375 

sulphate 377 

sulphide     ....!.  377 

white 376 

Le  Blanc  soda  process  .     .     .  282 

Levulose 415 

Lime 302 

air-slaked 303 

hypochlorite 307 

kilns 303 

slaked 303 

Lime  light 38 

Limestone «,.  305 

Limewater 303 

Limonite 339 

Litharge 375 

Lithium  V 294 

Luminosity  of  flames      .     .     .  216 

Lunar  caustic 366 

Magnesia 318 

alba 319 

usta 318 

Magnesite 318 

Magnesium '.     .  317 

basic  carbonate  .     .     .     .  319 

carbonate 318 

cement 318 

chloride 319 

family 316 


PAGE 

Magnesium  hydroxide    .     .     .  318 

oxide 318 

silicates .  321 

sulphate 320 

Magnetite 339,  349 

Malachite 357 

Manganates 381 

Manganese 379 

a  base-forming  element    .  380 

an  acid-forming  element  .  381 

in  glass 263 

occurrence 379 

oxides 380 

Manganic  acid 381 

Manganous  salts 380 

Marble 305 

Marl 305 

Marsh  gas 202 

Matches 242 

Matte 358 

Matter,  classification  ....  6 

conservation 5 

definition 5 

kinds 9 

Measurement  of  gases    ...  23 

Mechanical  mixtures       ...  6 

Meerschaum 321,336 

Mercuric  chloride 363 

iodide 364 

oxide 14,  362 

sulphide 363 

Mercurous  chloride    ....  363 

Mercury  ........  361 

iodides  .     .     .     ...     .  364 

metallurgy 361 

occurrence 361 

oxides 362 

uses 362 

Metaboric  acid .     .     .     .     .  -  .  265 

Metallurgy 268 

Metals 165,  267 

action  on  salts    ....  271 

definition 267 

extraction 268 

occurrence 267 

preparation  of  compounds  269 

reduction  from  ores     .     .  268 

Metaphosphoric  acid      .     .     .  245 

Metarsenic  acid 250 

Metasilicic  acid 261 

Metastannic  acid 371 


INDEX 


429 


PAGE 

Methane 202,  399 

Methylamine 412 

Mexican  onyx 305 

Mica 261,  336 

Microcosmic  salt 244 

Milk 414 

Minerals 267 

Minium 375 

Mixed  salts 244 

Molasses .  413 

Molecular  weights      ....  226 

boiling-point  method  .     .  233 

compared  with  oxygen     .  228 

determination      ....  2l5 

freezing-point  method       .  233 

oxygen  standard      .     .     .  227 

of  elements 232 

vapor-density  method       .  229 

Molecule 62 

Molybdenum 388 

Molybdic  acid 388 

Monazite  sand 377 

Mordants 333 

Morphine 418 

Mortar 304 

Moth  balls 418 

Muriatic  acid 182 

Naphthalene 418 

Naphthas 400 

Nascent  state 182 

Natural  gas 219 

sciences i 

Neon 80 

Neutralization 108 

a  definite  act 109 

definition 109 

heat  of  .     .     .     . •  .     .     .  109 

partial in 

Niagara  Falls 269,  329 

Nickel •  ....  354 

coin 359 

compounds 354 

plating 354 

Nicotine 419 

Nitrates 131 

Nitric  acid 128 

action  on  meta's      .     .     .  130 

decomposition     ....  129 

oxidizing  actirn  ....  130 
preparation          .     .     .  128,  140 


132 

78 


PAGE 

Nitric  acid  properties'     .     .    '.  129 

salts  ....'..'..  131 

Nitric  oxide 133 

Nitrites 132 

Nitrobenzene 417 

Nitrocellulose 415 

Nitrogen 78 

compounds 122 

in  air "8^" 

occurrence      .     .     .     .    78,  122 

oxides 

preparation     .... 

properties 80 

Nitroglycerin 409 

Nitrosulphuric  acid    ....  155 

Nitrous  acid 132 

oxide 132 

Non-metals 165 

Oil  of  myrbane 417 

of  vitriol 1 54 

Oleic  acid      .......  408 

Olein 409 

Oleomargarine 410 

Onyx 260 

Opal 260 

Open-hearth  process       .     .     .  346 

Opium 418 

Ores 267 

Organic  bases 412 

chemistry 201,397 

matter,  decomposition      .  122 

Orpiment 246 

Orthoarsenic  acid 250 

Orthophosphates 244 

Orthophosphoric  acid     .     .     .  244 

Orthosilicic  acid 261 

Osmic  acid 393 

Osmium    .    -.' 393 

tetroxide 393 

Oxalic  acid   ........  407 

Oxidation    •.*""". .     .     .     .     17,353 

definition 18 

Oxidizing  agent  <*r—T~ ...  37 

Oxygen 13 

and  ozone 22 

commercial  preparation    .  16 

history 13 

importance 21 

in  air,  estimation      ...  85 

in  air,  function    ....  84 


430     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


PAGE 

PACK 

Oxygen,  occurrence    .     .     . 

•        13 

Physical  changes  .     . 

...            2 

preparation     .... 

I3 

properties  .     .     . 

.     .    .        3 

properties  

16 

properties    and 

periodic 

standard   for  atomic 

groups    .     .     . 

...     171 

weights  

.      66 

state      .     .     .     . 

•     •    •        3 

two  atoms  in  molecule  . 

.     227 

Physics     

...     1,4 

Oxyhydrogen  blowpipe  .     . 

•       37 

Pitchblende.     .     .     . 

...     314 

Ozone  

21,137 

Plaster  of  Paris      .     . 

...     308 

Platinic  chloride    .     . 

•    •     •    393 

Palladium      ...... 

•     39° 

Platinized  asbestos    . 

•     •     •     391 

Palmitic  acid     

•     407 

Platinous  chloride 

•     •     •    393 

Palmitin  

•     409 

Platinum  

...     391 

Paraffin    

.     400 

a  catalytic  agent 

.     .152,392 

Paris  green  

.     250 

Pneumatic  trough 

.     .     .       16 

Parkes's  method  for  silver 

•     364 

Polyboric  acid  .     .     . 

...    265 

Pearls  

•     3°5 

Polyhalite     .     .     .     . 

...     288 

Perchloric  acid       .... 

.     187 

Polysilicic  acids     .     . 

...     261 

Periodic  acid     

•     194 

Porcelain      .     .     .     . 

•     •     •     336 

Periodic  division  .... 

.     166 

Portland  cement    .     . 

...    304 

groups  

.     167 

Potash      

...     293 

law    

.     169 

Potassium     .     .     .     . 

...     287 

law,  imperfections  .     . 

.     172 

acid  carbonate    . 

...    294 

law,  value  .     .     .     .     . 

.     171 

acid  sulphate 

...     294 

table      

.     168 

acid  sulphite  .     . 

...     294 

table,  arrangement  .     . 

.     166 

alum,  aluminium 

•     •     •     334 

Permanent  hardness  .     .     . 

.     310 

alum,  chrome 

...     384 

Permanganates      .... 

.     381 

alum,  iron  .     .     . 

•     •     •     352 

Permanganic  acid  .... 

.     381 

and  plant  life 

...     287 

Peroxides      

.     278 

aurate    .     .     .     . 

...     396 

Petroleum    

•     399 

bromide     .     .     . 

...     .     290 

Pewter     

•     372 

carbonate  .     .     . 

'  •<     •     •     293 

Phenol      

.     417 

chlorate     .     .     . 

.     .     .     291 

Philosopher's  stone   .     .     . 

9 

chloride     .     .     . 

...     290 

Phlogiston    

.       19 

chromate    .     .     . 

...     385 

Pfesphates  

•     245 

cyanide  .... 

.     .     .     293 

Phosphine    

.     242 

dichromate     .     . 

...     386 

Phosphonium  compounds  . 

•     243 

ferricyanide    .     . 

•     •     •     352 

Phosphoric  acid    .... 

•     244 

ferrocyanide   .     . 

•     •     •     352 

Phosphorite      

•     239 

hydroxide  .     .     . 

...     288 

Phosphorous  acid      .     .     . 

.     244 

hydroxide,  action 

of  halo- 

Phosphorus  

•     239 

gens   .     .     .     . 

.  N.     .     289 

acids      

•     243 

hypochlorite  .     . 

...     289 

family    

.     238 

iodide    .     .     .     . 

.     .     .     290 

hydrogen  compounds  . 

.     242 

manganate      .     . 

...     38i 

occurrence      .     .     . 

•     239 

nitrate    .     .     .     . 

...     291 

oxides    

•     243 

occurrence      .     . 

...     287 

preparation     .... 

•     239 

permanganrte 

...     381 

properties  

.     240 

preparation     .     . 

...     288 

red    

.     241 

sulphate     .     .     . 

.    ...     294 

yellow    

.     240 

Precipitated  chalk 

•     «     •     3°6 

Photography     

.     367 

Precipitation     .     .     . 

...     140 

INDEX 


431 


PAGE 

Properties,  chemical ....  3 

physical 3 

Prussic  acid 210 

Puddling 343 

furnace 344 

Pyridine 412 

Pyrites 351 

Pyrolusite 380 

Pyrophosphoric  acid ....  245 


Quantitative  equations  . 

Quartz 

Quicklime 

Quinine 

Quinoline 


Radical 

Radium 

Reaction,  classes 

addition 

completed 

heat  of 

of  decomposition    .     .     . 

of  double  decomposition . 

of  substitution    .     .     .     . 

reversible  . 


steps  in      ... 

Realgar 

Red  lead 

phosphorus     .     . 
Reducing  agent 
Reduction    .     .     .     . 

Rennin 

Resemblances,  family 
Respiration  .... 
Rhodium  .... 
Rochelle  salts  .  .  . 
Rouge  .'  .  .  .  . 
Rubidium  .... 

Ruby 

Ruthenium  .... 
Rutile  . 


36, 


72 
260 
302 
418 
412 

112 

313 

70 

70 

139 

75 
70 

7i 
70 

137 

I3< 

246 

375 
241 

37 
354 
414 
170 

87 
390 
408 

349 
294 

33i 

390 
264 


Safety  lamp 202 


Sal  ammoniac 

soda. 
Salt  .  . 
Saltpeter . 

Chili 
Salts  .  . 

acid  . 


294 
282 
280 
291 
285 
109 

112 


PAGE 

Salts  basic 1 1 1 

binary 114 

characteristics     ....  109 

definition 109 

insoluble 272 

mixed 244 

nomenclature      .     .     .     .  113 

normal 112 

preparation  by  precipita- 
tion      270 

Sand 260 

Sandstone 260 

Saponification 410 

Sapphire 331 

Satinspar 308 

Scale 320 

Schonite 288 

Selenite 308 

Selenium 161 

Serpentine 320,  336 

Shot 247,375 

Siderite 339 

Silica 260 

Silicates  ........  261 

Silicic  acids 261 

Silicides 259 

Silicon 258 

acids 261 

dioxide .......  260 

fluoride 258 

hydride 258 

Silver  .     .     .     .     .     .     .     .     .  364 

amalgamation  process      .  364 

bromide 367 

chloride 367 

coin 359 

German 359 

in  copper  ores     ....  358 

iodide  * 367 

metallurgy 364 

nitrate 366 

oxide     . 366 

parting  of 365 

refining 365 

sulphide 366 


Slag 

Smalt  ...... 

Smithsonite .  .  . 
Smokeless  powder 
Soaps 


.  .  340 

•'  •  355 

.  .  321 

.  .  293 

.  .  410 

Soda  ash 284 


432     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


PAGE 

PAGE 

Soda  lime     

.       .       202 

Stannous  chloride      .     .     .     . 

372 

Sodium    

.     .     276 

Starch       

4J5 

acetate  

.     .     407 

Stassf  urt  salts  -. 

287 

bicarbonate    .     .     . 

.     .     285 

Stearic  acid  

407 

carbonate  .... 

.     .     282 

Stearin     

409 

carbonate,  historical 

.     .     284 

Steel    

345 

chloride      .... 
chromates  .... 

.     .     280 
.     .     386 

alloys     .  \    
properties  ,\  

348 
347 

hydrogen  carbonate 

.     .     285 

tempering  of  

348 

hydroxide  .... 

.     .     278 

tool  

347 

hyposulphite  .     .     . 

.     .     282 

Stibine      

251 

iodate    

.     .     191 

Stibnite    

250 

nitrate   

.     .     285 

Stoneware     

336 

occurrence      .     .     . 

.     .     276 

Strontianite  

312 

peroxide     .... 

.     .     277 

Strontium     

312 

phosphates      .     .     . 

.     .     286 

hydroxide  1     

312 

preparation 

.     .     276 

nitrate    .    v    

312 

properties  .... 

.     .     277 

Structural  formulas    .... 

119 

sulphate      .... 

.     .     281 

Structure  of  compounds     .     . 

119 

sulphite      .... 

.     .     281 

Strychnine    

419 

tetraborate      .     .     . 

.     .     287 

Substitution  

70 

thiosulphate   .     .     . 

.     .     282 

Sugars      

412 

Solder      

•  372,375 

cane  .     .  »  

412 

Solubility  of  gases      .     . 

•     .       95 

fruit  

415 

of  solids     .... 

.     .       96 

grape      

414 

Solution   

.     .       94 

milk       

414 

and  chemical  action 

•     •       53 

Sulphates      

iSQ 

boiling  point  .     .     . 

.     .       98 

Sulphides      

148 

classes    

.     .       94 

Sulphites       

IS2 

distribution  of  solids 

in    .       98 

action  of  acids  on    ... 

ISO 

electrolysis  of      .     . 

.     .       99 

Sulphur    

M3 

freezing  point      .     . 

.     .       99 

allotropic  forms       .     .     . 

144 

of  gases  in  liquids  . 

.     .       94 

chemical  properties      .     . 

145 

of  solids  in  liquids  . 

.     .       96 

comparison  with  oxygen  . 

161 

properties  .... 

.     .       98 

dioxide  

149 

saturated    .... 

.     .       97 

preparation  .... 

149 

supersaturated    .     . 

.     .       98 

properties    .... 

'•5° 

Solvay  soda  process  .     . 

.     .     283 

extraction  

M3 

Sombrerite    

.     .     239 

flowers  of  

143 

Spectroscope      .... 

.     .     296 

occurrence      

M3 

Sphalerite     

•     •     325 

oxides    

149 

Spiegel  iron       .... 

•     •     343 

physical  properties       .     . 

144 

Spinel       

•     •     332 

trioxide       

152 

Spontaneous  combustion 

.     .       20 

uses  . 

146 

Stalactites     

••    •     3°5 

varieties     

144 

Stalagmites  

•    •     305 

Sulphuric  acid  

"54 

Standard  conditions  .     . 

•     •      23 

action  as  an  acid      .     .     . 

iS7 

Stannates      

•     •     372 

action  on  metals      .     .     . 

T57 

Stannic  acid      .... 

•     •    372 

action  on  organic  matter 

158 

chloride      .... 

•     •     372 

action  on  salts    .... 

158 

oxide     

•     •    372 

action  on  water  .... 

158 

INDEX 


433 


PAGE 

Sulphuric  acid  fuming    .     .     .  155 

manufacture 154 

oxidizing  action       .     .     .  157 

plant 156 

properties 15? 

salts 159 

Sulphuric  anhydride  ....  153 

Sulphurous  acid 151 

Superphosphate  of  lime      .     .  246 

Sylvine 288 

Symbols n 

Synthesis 40 

Table,  alkali  metals  ....  274 
alkaline-earth  metals  .  .  300 
alloys  of  copper .  .  .  .  359 
aqueous  tension  Appendix  B 
atomic  weights  .  Appendix  A 
chlorine  family  ....  174 
composition  of  earth's 

crust  ..'.....  10 
composition  of  fuel  gases  220 
constants  of  elements 

Appendix  B 

copper  family  ....  356 
elements  .  .  .  Appendix  A 
gold  and  platinum  metals  390 
hydrocarbons  ....  399 
magnesium  family  .  .  .  316 
manganese  and  chromium  379 
periodic  arrangement  .  .168 
phosphorus  family  .  .  .  238 
silicon  family  .  .  .  .  257 
solubility  of  gases  in  water  95 
solubility  of  salts  ...  96 
solubility  of  salts  at  dif- 
ferent temperatures  .  97 

tin  and  lead 370 

weights  of  gases      Appendix  B 

Talc 321,336 

Tartar  emetic 408 

Tartaric  acid 408 

Tellurium 161 

Temporary  hardness       .     .     .     309 
Ternary  acids    .     .     .     .     .     .     113  . 

salts 114 

Tetraboric  acid      .     .     ...     .     265 

Thallium 327 

Theory,  atomic 61 

definition 64 

value  of 64 


PAGE 

Thermite 331 

Thio  compounds 282 

Thiosulphates 159 

Thiosulphuric  acid    ....  159 

Thorium 377 

Tin 370 

block 371 

compounds 372 

crystals 372 

family 370 

foil 371 

metallurgy 370 

plate 371 

properties 371 

uses 371 

Titanium 257,  264 

Topaz 331 

Triad  families 166 

Tungsten 388 

Type  metal  .     .     .     .     .     .253,375 

Uranium 388 

Valence 116 

a  numerical  property  .     .  116 

and  combining  ratios  .     .  118 

and  equations     ....  120 

and  formulas 120 

and  periodic  groups     .     .  162 

and  structure       .     .     .     .  118 

definition 116 

indirectly  determined  .     .  117 

measure  of 117 

variable 117 

Vaseline  .  • 400 

Venetian  red 349 

Verdigris 407 

Vermilion 363 

Vinegar    . 406 

Vitriol,  blue 361 

green 350 

oil  of 154 

white 324 

Volume  and  aqueous  tension  .  25 

and  pressure 24 

and  temperature      ...  23 

of  combining  gases      .     .  194 

Water 40 

a  compound 40 

and  disease 49 


434 


AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


PAGE 

Water,  catalytic  action  of  .     .  154 

chalybeate 351 

chemical  properties      .     .  53 

composition 47 

composition  by  volume    .  44 

composition  by  weight     .  47 

dissociation  of    ....  210 

distillation  of      ....  50 
electrolysis  of     .     .     .    41,  103 

filtration  of 51 

gas 219 

hard 309       Xenon 

historical 40 

impurities  in 48 

in  air 87 

mineral 49 

occurrence 48 

of  crystallization     .     .      54, 75 

physical  properties'.     .     .  53 

purification  of     ....  50 

qualitative  analysis       .     .  41 

quantitative  analysis    .     .  42 

river 49 

sanitary  analysis      ...  50 

self-purification  ....  53 

softening  of 310 

standard  substance      .     .  55 


PAGE 
Water,  synthesis 43 

uses  of  .......       55 

Weights,  atomic    .....       65 

Welsbach  mantles  .  .  .219,377 

Whisky 404 

Wine 404 

Witherite 312 

Wood  alcohol 402 

distillation 402 

Wood's  metal  ......  254 


80 


Yeast 403 

Zinc 321 

alloys  of 323 

-  blende 321 

chloride 325 

flowers  of 322 

metallurgy 321 

occurrence 321 

oxide 324 

sulphate 324 

sulphide 325 

white 324 

Zymase 403 


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